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MANUAL  OF  QUALITATIVE 
CHEMICAL  ANALYSIS 


BY 


J.  F.  McGREGORY 

\> 

PROFESSOR  OF  CHEMISTRY  AND  MINERALOGY  IN 
COLGATE  UNIVERSITY 


GINN   &  COMPANY 

BOSTON  •  NEW  YORK  •  CHICAGO  •  LONDON 


COPYRIGHT,  1903 
BY  J.  F.  McGREGORY 


ALL.  RIGHTS  RESERVED 
47.9 


J&rte* 


GINN   &   COMPANY-  PRO- 
PRIETORS •  BOSTON  •  U.S.A. 


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PREFACE 

An  examination  of  most  of  the  works  on  the  subject  of  quali- 
tative analysis  will  show  that  they  belong  in  one  of  two  general 
classes  which  may  be  described  as  follows  :  first,  the  exhaustive 
treatise,  of  which  such  a  work  as  that  of  Fresenius  will  serve 
as  an  example  ;  second,  the  abbreviated  treatise,  often  very 
much  abbreviated,  in  which  the  author  attempts  to  cover  the 
whole  work  in  a  few  lessons. 

For  graduate  students,  or  for  the  better  class  of  beginners 
who  are  able  to  devote  the  most  of  their  time  to  the  study,  books 
of  the  first  class  may  be  invaluable.  But  the  great  majority  of 
the  students  in  our  colleges  do  not  become  chemists.  They 
study  analytical  chemistry  at  the  most  but  a  short  time,  and 
what  they  acquire  is,  and  should  be,  to  a  considerable  extent,  of 
disciplinary  value  to  them.  To  the  great  majority  of  our  stu- 
dents, therefore,  the  exhaustive  treatise,  especially  at  the  begin- 
ning of  their  course,  is  a  means  of  confusion  rather  than  an 
intelligent  guide. 

The  second  class  of  text-books  is  also  likely  to  fail  with  the 
average  student,  since  it  teaches  him  to  analyze  an  unknown 
substance  in  such  a  purely  mechanical  way  that  the  actual  knowl- 
edge of  the  subject  acquired  is  small,  and  the  disciplinary  value 
of  the  work  becomes  a  minimum. 

This  little  manual  was  first  written  to  meet  the  wants  of  the 
author's  own  classes,  in  which  it  has  been  used  for  several  years, 


M44288 


iv  PREFACE 

and  has  now  been  carefully  revised  for  this  edition.  In  it  the 
attempt  has  been  made  to  retain  the  essentials  of  the  larger 
works,  omitting  the  rare  metals  and  acids,  and,  at  the  same  time, 
to  avoid  the  "short  cuts"  so  often  found  in  smaller  works.  It 
is  presupposed  that  the  student  has  had  a  thorough  course  of 
instruction  in  general  chemistry,  at  least  through  the  non- 
metallic  elements,  before  beginning  this  work. 

All  laboratory  work  ought  to  be  carried  on  under  the  imme- 
diate supervision  of  a  competent  instructor.  The  work  should 
also  be  accompanied  by  a  sufficient  number  of  examinations  to 
bring  out  all  the  essential  points  connected  with  the  work. 
The  author's  own  plan  is  to  give  frequent  oral  examinations 
throughout  the  course,  especially  in  the  earlier  parts. 

In  the  introduction  will  be  found  certain  definitions  and 
general  principles,  which  the  student  should  know  at  the  begin- 
ning of  this  course.  The  author  has  thought  it  best  to  omit 
all  consideration  of  the  dissociation  theory,  believing  that,  how- 
ever valuable  the  study  of  this  subject  may  be  to  the  chemist, 
its  introduction  as  a  basis  of  study  in  qualitative  analysis  is  not 
to  be  recommended,  and  that  its  consideration  should,  therefore, 
be  deferred  until  a  later  time  when  the  student  shall  have  a 
larger  number  of  facts  at  his  command. 

Special  attention  ought  to  be  given  to  Parts  I  and  II,  which 
deal  with  simple  substances  only.  The  reactions  here  given  are 
often  partially  or  entirely  omitted  in  a  text-book  ;  but  since 
they  form  the  basis  of  all  the  more  advanced  portions  of  the 
work,  they  ought  to  be  thoroughly  mastered. 

In  the  separation  of  the  metals  in  Part  III  only  one  prac- 
tical and  well-established  method  is  given,  it  being  the  author's 


PREFACE  V 

experience  where  several  methods  are  given,  either  that  only 
one  is  used  or  that  the  student  is  likely  to  get  them  confused. 
Later  in  the  course  the  student  may  learn  other  methods  as  it 
seems  desirable. 

Part  IV  has  been  inserted  in  order  to  make  the  work  more 
complete.  It  is  not  essential  for  all  students,  and,  if  it  is  found 
necessary  to  shorten  the  course,  may  be  omitted. 

The  appendix  contains  tables  and  some  useful  information 
for  both  students  and  instructor. 

The  author  desires  to  express  his  thanks  to  his  assistant, 
Mr.  R.  B.  Smith,  and  to  all  others  who,  either  by  suggestion  or 
criticism,  have  so  kindly  assisted  him  in  this  work ;  also  to  Pro- 
fessor R.  W.  Thomas  for  his  careful  reading  and  criticism  of 

the  manuscript. 

J.  F.  M. 
HAMILTON,  N.Y.,  September,  1,  1903. 


CONTENTS 

PAGE 
INTRODUCTION    .  ...  xi 

PART  I 

REACTIONS  FOR  THE  METALS  IN  SOLUTION  1 

.  Lead .1 

Silver 3 

i  Mercury  (Mercurous)   .........       4 

'  Mercury  (Mercuric) 6 


Bismuth 
Copper                                   ....... 

.       7 
8 

9 

Arsenic 

10 

Antimony     ... 
Tin  (Stannous)          .... 
Tin  (Stannic)       ....                  . 
Aluminum         ...... 
Chromium    .......... 

.     10 
11 
.     12 
13 
.     14 

Iron  (Ferrous) 
Iron  (Ferric) 
Nickel       

15 
.     16 
17 

Cobalt           .         .                  ... 

.     19 

Manganese 
Zinc      

20 
.     22 

Magnesium       ......                  . 
Barium          

23 
.     24 

Strontium          

25 

Calcium       ".***&•*«'•.'•''' 

- 

Potassium 

.     26 

27 

Sodium         

.     28 

•.Ammonium       .                             . 

29 

REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION 

.     30 

Hydrochloric  Acid    ........ 
Hydrobromic  Acid        ........ 
vii 

31 
.     31 

viii  CONTENTS 

REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION  —  Continued     PAGE 

Hydriodic  Acid          .....                  ...  32 

Hydrofluoric  Acid         ...                  32 

Hydrocyanic  Acid 33 

Sulfocyanic  or  Thiocyanic  Acid 34 

Hydroferrocyanic  Acid 34 

Hydroferricyanic  Acid 34 

Hypochlorous  Acid  .........  35 

Chloric  Acid         ..........  35 

Hydrogen  Sulfid  (Hydrosulfuric  Acid)      .          .          .          .         .  36 

Thiosulfuric  Acid 36 

Sulfurous  Acid          .........  37 

Sulfuric  Acid 38 

Chromic  Acid  .         .                   38 

Nitrous  Acid         ......                  39 

Nitric  Acid       .                  .                  40 

Phosphoric  Acid  .                   41 

Arsenious  Acid                    .          .                   41 

Arsenic  Acid         .          .          .          .  -  .         .          .          .          .42 

Boric  Acid         .                  .                  43 

Carbonic  Acid       ..........  43 

Silicic  Acid       ..........  44 

Acetic  Acid  .         .         .         .         .         .         .         .         .         .         .  45 

Oxalic  Acid       ..........  46 

Tartaric  Acid        ..........  46 

PART  II 

REACTIONS   FOR   DRY   SUBSTANCES 

BLOWPIPE  ANALYSIS       .         . 48 

The  Effect  of  Heat  alone 50 

The  Substance  is  heated  on  Charcoal         .....  55 

The  Substance  is  heated  on  Charcoal  with  Sodium  Carbonate      .  58 

Coloration  of  the  Flame    .         .         .         .                           .         .  59 

Coloration  of  the  Borax  or  Microcosmic  Bead       .         .         .         .60 

The  Substance  is  fused  on  Platinum  Foil  with  Sodium  Carbonate 

and  Potassium  Nitrate       .......  62 

The  Substance  is  acted  upon  by  Sulfuric  Acid      .         .         .         .62 

Special  Tests 66 


CONTENTS  ix 

PART  III 

SYSTEMATIC   EXAMINATION   FOR   METALS   IN   SOLUTION 

PAGE 

SIMPLE  COMPOUNDS 69 

Group  1    .  .  .  69 

Group  2         ...  70 

Group  3    ...  72 

Group  4 .     73 

Group  5 73 

Group  6        .*....'.  .  .74 

Examination  for  Acid  Radicals         ......         75 

SYSTEMATIC   EXAMINATION  FOR   METALS   IN   SOLUTION 

MIXED  COMPOUNDS 76 

Preliminary  Examination 77 

Group  1  —  Lead,  Silver,  Mercury  (Mercurous)      .          .         .          .78 
Group  2 — Mercury  (Mercuric),  Bismuth,  Copper,  Cadmium,  Ar- 
senic, Antimony,  Tin          .......  80 

Group  2,  Subdivision  A 82 

Group  2,  Subdivision  B 84 

Group  3  —  Aluminum,  Chromium,  Iron 86 

Phosphates,  Oxalates,  etc.,  are  absent        ....  88 

Phosphates,  Oxalates,  etc.,  are  present 90 

Group  4  —  Nickel,  Cobalt,  Manganese,  Zinc      ....  93 

GroupS  —  Barium,  Strontium,  Calcium 96 

Group  6  —  Magnesium,  Potassium,  Sodium,  Ammonium           .  98 

SYSTEMATIC   EXAMINATION  FOR  ACID   RADICALS   IN   SOLUTION 

Preliminary  Examination  ......  .101 

Preparation  of  the  Solution     .         .          .          .          .          .          .  102 

Classification  of  the  Acid  Radicals      ....  103 

Acids  :  Group  1  104 

Acids :  Group  2 .           .  107 

Acids:  Group  3       .                                                ....  110 


x  CONTENTS 

PATTT  IV 

PAGE 

SYSTEMATIC  EXAMINATION  OF  COMPLEX   SOLIDS           .          .111 
Preliminary  Examination 112 

I.  THE  SUBSTANCE  is  A  METAL  OK  AN  ALLOY     ....     113 

A.  Metals  insoluble  and  unchanged  in  nitric  acid  .         .         .         113 

B.  Metals  which  form  insoluble  oxids  by  the  action  of  nitric 

acid 114 

C.  Metals  and  alloys  soluble  in  nitric  acid      .         .         .         .  114 

• 

II.  THE  SUBSTANCE  is  NEITHER  A  METAL  NOR  AN  ALLOY  .         .  115 

A.  The  substance  is  partially  or  entirely  soluble  in  water        .         116 

B.  The  substance  is  insoluble  in  water        .         .          .         .          .116 

C.  The  substance  is  insoluble  in  H2O  and  in  HC1  .         .  117 

D.  The  substance  is  insoluble  in  H2O  and  in  both  HC1  and  HNO3  1 17 

E.  The  substance  is  insoluble  in  H2O  and  in  all  acids    .         .  118 

F.  The  substance  is  a  silicate 120 

(a)  Silicates  decomposed  by  acids 121 

(b)  Silicates  not  decomposed  by  acids          .         .         .         .121 

G.  Cyanids  are  present 122 

APPENDIX 

NAMES,  SYMBOLS,  AND  ATOMIC  WEIGHTS  OF  THE  ELEMENTS        .  125 

NAMES  AND  FORMULAS  OF  REAGENTS  AND  SOLUTIONS     .         .  126 

PREPARATION  OF  REAGENTS  AND  SOLUTIONS          ....  129 

INDEX   .......  131 


INTRODUCTION 

Analytical  Chemistry  treats  of  the  composition  of  substances 
and  of  the  methods  by  which  we  determine  the  same.  There 
are  two  general  divisions  of  analytical  chemistry,  viz. :  qualita- 
tive and  quantitative  analysis. 

Qualitative  Analysis  has  for  its  object  the  determination  of 
the  constituent  elements  of  a  body.  It  consists  in  the  separa- 
tion of  each  of  the  elements,  either  in  the  free  state,  or,  as 
more  commonly  happens,  in  the  form  of  some  compound  which 
is  characteristic  and  easily  recognized. 

Quantitative  Analysis  belongs  to  a  more  advanced  course, 
and  has  for  its  object  the  determination  of  the  percentage 
amounts  of  the  constituents  of  a  body,  and  thus  of  the  actual 
constitution  of  the  body. 

Every  simple  inorganic  substance  consists  of  two  parts.  The 
first,  which  is  a  metal  or  positive  radical,  is  chemically  com- 
bined with  the  second,  which  is  a  non-metal  or  negative  radical. 
The  more  complex  substances  may  contain  several  metals,  or 
positive  radicals,  and  often  contain  more  than  one  acid,  or  nega- 
tive radical.  These  may  be  chemical  combinations  or  merely 
mechanical  mixtures. 

By  subjecting  a  substance  to  various  conditions  we  obtain  a 
series  of  phenomena  which  we  call  its  reactions  ;  and  any 
known  substance  which  is  employed  in  effecting  a  reaction  is 
called  a  reagent.  The  subjecting  of  a  substance  to  the  action 
of  reagents,  by  means  of  which  its  constituent  elements  are 
recognized,  is  the  process  employed  in  qualitative  analysis. 

We  may  subject  the  substance  to  the  action  of  reagents 
either  in  its  original  solid  condition  —  if  it  be  a  solid  —  or  in 

xi 


xii  INTRODUCTION 

solution.  These  two  methods  of  examination  are  known  as  the 
dry  way  and  the  wet  way. 

The  dry  way  may  be  employed  for  the  complete  analysis  of 
simple  substances,  and  is  a  valuable  aid  in  making  preliminary 
tests  of  complex  substances.  This  comprises  what  is  known 
as  Blowpipe  Analysis,  which  is  fully  explained  in  Part  II  of 
this  work. 

The  wet  way  is  more  generally  used  in  qualitative  analysis, 
because  its  reactions  are,  for  the  most  part,  simpler  and  more 
rapid.  It  can  be  employed  with  all  kinds  of  substances.  Solids 
and  gases  can  generally  be  obtained  in  solution  in  water,  or  some 
other  convenient  liquid,  in  which  they  will  dissolve  without 
losing  their  characteristic  properties. 

When  a  substance  in  solution  is  acted  upon  by  a  reagent  the 
results  are  always  in  accordance  with  a  law  which  may  be 
stated  as  follows:  When  two  substances  which  are  in  contact  in 
solution  can,  under  the  conditions  of  the  reaction,  form  a  sub- 
stance which  is  insoluble  or  volatile,  the  insoluble  or  volatile 
substance  will  always  be  formed  and  continue  to  be  formed  until 
one  of  the  factors  is  exhausted.  An  insoluble  compound  thus 
formed  is  called  a  precipitate,  and  precipitation  is  the  most 
common  form  of  reaction  in  analytical  chemistry. 

A  precipitate  may  be  of  almost  any  color  and  the  color  may 
be  characteristic,  or  if,  as  is  more  often  the  case,  it  is  not 
especially  so,  other  precipitates  are  formed  with  other  reagents 
until  the  combination  of  results  is  such  as  to  determine  the 
substance  with  certainty.  Sometimes  a  precipitate  which  is 
not  characteristic  becomes  so  by  its  solubility  or  insolubility 
in  some  other  reagent  or  in  an  excess  of  the  reagent  first  used. 
Solubility  in  excess  is  confined  to  a  few  reagents.  These  are 
the  alkaline  hydroxids  and  a  few  alkaline  salts. 

The  treatment  of  a  substance  with  a  reagent  sometimes 
results  in  the  formation  of  a  gas,  which  is  recognized  by  its 
color  or  odor  or  by  some  other  characteristic  property.  This 


INTRODUCTION  xiii 

almost  always  results  from  the  decomposition  of  some  acid 
radical  by  means  of  some  acid  used  as  a  reagent. 

In  order  to  avoid  constant  mistakes,  the  student  should 
thoroughly  understand  every  reaction  which  he  uses.  In 
acquiring  this  knowledge  he  should,  especially  in  the  earlier 
and  simpler  portion  of  the  course,  accustom  himself  to  the  use 
of  all  the  common  reagents.  In  the  more  advanced  portion  of 
his  course  a  more  systematic  and  selective  use  of  reagents  will 
be  necessary. 

The  student  should  first  be  required  to  perform  all  the  reactions 
in  Part  I.  He  should  write  out  the  chemical  equation  in  every 
case  in  a  suitable  notebook,  and  at  the  time  the  reaction  is  made. 
He  may  then  be  given  some  simple  solutions  and,  by  the  use  of 
the  reactions  he  has  just  been  performing,  find  what  the  unknown 
solution  contains.  Having  found  the  substance  contained  in  a 
solution,  he  should  try  all  the  reactions  given  for  that  substance. 
He  should  also  compare  each  reaction  with  all  similar  reactions 
given  by  other  substances,  noting  points  of  difference,  and  in 
this  way  make  each  reaction  as  comprehensive  as  possible. 

Many  students  seem  to  feel  that  all  that  is  required  of  them 
is  to  find  out  what  is  in  the  unknown  solution.  This  is 
undoubtedly  the  goal  toward  which  their  course  of  study  is 
tending  ;  but  if,  at  the  beginning  of  his  course,  that  is  all  that 
the  student  desires,  a  much  shorter  method  would  be  to  ask 
the  instructor.  The  value  of  an  elementary  course  of  study  in 
analytical  chemistry  is  not  simply  to  find  out  what  a  solution 
or  solid  substance  contains,  but  to  learn  how  to  find  out  what 
it  contains. 

It  will  be  observed  that  in  Part  III  all  tables  such  as  are 
often  included  in  a  text-book  and  intended  for  aid  in  the  sepa- 
ration of  the  metals  have  been  omitted.  This  is  because,  in 
the  author's  judgment,  the  use  of  them  in  the  hands  of  the 
majority  of  students  is  pernicious.  When  such  tables  are  used 
the  student  almost  always  depends  upon  them  rather  than  upon 


xiv  INTRODUCTION 

the  full  text,  and  so  is  frequently  led  into  error  because  of  some- 
thing which  has  been  omitted.  The  information  given  in  a 
table  is  necessarily  brief,  and  the  average  student  can  acquire 
such  information  by  a  few  hours  of  study.  With  such  informa- 
tion in  his  head  he  will  work  much  faster  and  with  more  satis- 
faction to  himself  and  his  instructors,  and  when  in  doubt  he 
will  consult  the  text  and  not  a  table. 

Every  student  should  be  required  to  keep  a  notebook  in 
which  to  record  the  results  of  all  his  work  in  the  laboratory. 
He  should  also  be  encouraged,  if  not  required,  to  make  use  of 
some  large  and  fairly  complete  text-book  on  general  chemistry 
for  collateral  reading,  especially  in  connection  with  those  com- 
pounds which  he  meets  with  in  the  course  of  his  work.  Such  a 
course  of  action,  if  persisted  in,  will  give  to  the  student  a  much 
more  comprehensive  view  of  the  subject,  and  will,  in  addition, 
provide  him  with  a  fund  of  information  which  will  always  be  of 
value  to  him. 

The  instructor  should  require  the  student  to  do  clean  and 
careful  work.  Work  done  in  a  careless  way,  with  dirty  appa- 
ratus and  on  a  dirty  desk,  is  of  little  or  no  value.  Clean,  intel- 
ligent work,  accompanied  by  a  reasonable  amount  of  reading 
and  study,  will  give  to  the  student,  even  if  his  course  is  only  a 
short  one,  a  glimpse  at  least  of  the  immense  and  interesting 
field  of  study  and  research  which  is  always  open  to  the  chemist 


QUALITATIVE    ANALYSIS 

PART  I 

REACTIONS  FOR  THE  METALS  IN  SOLUTION 

LEAD,  Pb" 

Lead  dissolves  easily  in  HNO3  with  formation  of  Pb(NO3)2. 

3  Pb  +  8  HN03  =  3  Pb(N08)2  +  4  H20  +  2  NO. 
It  dissolves  in  hot  concentrated  H2SO4. 

Pb  +  2  H2S04  =  PbS04  -f  2  H20  +  S02. 

It  is  not  attacked  by  dilute  H2SO4  or  HC1. 
For  the  reactions  use  lead  nitrate,  Pb(NO3)2. 

1.  Sodium  Hydroxid  precipitates  white  Pb(OH)2  or  a  white  basic 
hydroxid,  Pb2O(OH)2,  according  to  the  conditions  which  exist. 

Pb(N03)2  +  2  NaOH  =  Pb(OH)2  +  2  NaN03. 
2  Pb(N03)2  +  4  NaOH  =  Pb20(OH)2  +  4  NaNO3  +  H20. 

Soluble  in  excess  of  the  reagent  (4  vols.),  easily  soluble  in  con- 
centrated NaOH,  forming  sodium  plumbite,  Na2PbO2. 

Pb(OH)2  +  2  NaOH  =  Na2PbO2  +  2  H20. 

2.  Ammonium     Hydroxid     precipitates     a     white    basic     salt, 
(PbO)2Pb(N03)2. 

3  Pb(N08)2  -1-  4  NH4OH  =  (PbO)2Pb(N03)2  +  4  NH4N08  -f  2  H20. 

Insoluble  in  excess  of  the  reagent. 

1 


2  ...    .  .    QUALITATIVE   ANALYSIS 

3.  .Sodium  /xr  Ammonium.  Carbonate  precipitates  white  PbCO3. 

If  the  solution  is  hot  a  white  basic  carbonate,  Pb3(OH)2(CO3)2, 
is  precipitated.     This  is  known  commercially  as  white  lead. 

3Pb(N03)2  +  3Na2C03  +  H20  =  Pb3  (OH)2(C03)2  +  6NaN03  +  C02. 

4.  Hydrogen  or  Ammonium  Sulfid  precipitates  black  PbS. 

Pb(N03)2  +  (NH4)2S  =  PbS  +  2  NH4N08. 

Insoluble  in  cold  dilute  acids.     Soluble  in  warm  dilute  HNO3, 
forming  Pb(NO3)2  and  free  sulfur. 

3  PbS  +  8  HN03  =  3  Pb(N03)2  +  4  H20  -f  2  NO  +  3  S? 
Concentrated  HNO3  oxidizes  PbS  to  PbSO4. 

3  PbS  +  8  HN03  =  3  PbS04  +  4H20  +  8  NO. 

If  the  HNO3  is  of  medium  strength  both  reactions  will  go  on 
at  the  same  time. 

5.  Acid  Sodium  Phosphate  precipitates  white  Pb3(PO4)2. 

3  Pb(N03)2  +  2  Na2HP04  =  Pb3(P04)2  +  4  NaN03  +  2  HN03. 
Easily  soluble  in  HNO3. 

6.  Potassium  Cyanid  precipitates  white  Pb(CN)2. 

Pb(N03)2  +  2  KCN  =  Pb(CN)2  +  2  KN03. 
Insoluble  in  excess  of  the  reagent. 

7.  Potassium  lodid  precipitates  yellow  PbI2. 

Pb(N03)2  +  2  KI  =  PbI2  +  2  KN03. 

The  precipitate  is  soluble  in  boiling  water  (4  vols.),  from  which 
solution  it  crystallizes,  on  cooling,  in  golden  yellow  scales. 

8.  Potassium  Chromate  precipitates  yellow  PbCrO4. 

Pb(N03)2  +  K2Cr04  =  PbCr04  +  2  KN03. 

Soluble  in  NaOH  (5  vols.  of  dilute  or  1  vol.  of  concentrated) 
and  in  HNO3.     Insoluble  in  acetic  acid. 


REACTIONS  FOR  THE  METALS  IN  SOLUTION  3 

9.    Potassium  Ferrocyanid  precipitates  white  Pb2Fe(CN)6. 
2  Pb(N03)2  +  K4Fe(CN)6  =  Pb2Fe(CN)6  +  4  KN03. 

10.  Hydrochlorid  Acid,  or  any  soluble  chlorid,  precipitates  white 

PbCl2. 

Pb(N03)2  +  2  HC1  =  PbCl2  +  2  HN03. 

Easily  soluble  in  boiling  water,  from  which,  unless  the  solution 
is  too  dilute,  it  will  crystallize,  on  cooling,  in  long  white  needles. 

11.  Sulfuric    Acid,    or   any    soluble    sulfate,     precipitates   white 
PbSO4. 

Pb(N03)2  4-  H2S04  =  PbS04  +  2  HN03. 

Easily  soluble   in   ammonium  tartrate  or  ammonium    acetate. 
[Add  tartaric  or  acetic  acid,  and  then  excess  of  NH4OH.] 

12.  Metallic  Zinc  will  entirely  precipitate  the  lead  in  crystalline 
form. 

Pb(N03)2  +  Zn  =  Pb  +  Zn(N03)2. 

NOTE.  The  chemical  equations  have  been  given  under  Lead  as  a  guide 
for  the  student.  Such  equations  will  generally  be  omitted,  but  the  student 
should  be  required  to  write  them  for  himself. 

SILVER,  Ag' 

Silver  dissolves  easily  in  HNO3  with  formation  of  AgNO3. 
Insoluble  in  HC1  and  in  H2SO4. 

For  the  reactions  use  silver  nitrate,  AgNO3. 

1.  Sodium  Hydroxid  precipitates  brown  Ag2O.     Insoluble  in 
excess.     Soluble  in  NH4OH,  forming  NH4AgO. 

2.  Ammonium  Hydroxid  precipitates  the   same.     Very  easily 
soluble  in  excess,  forming  NH4AgO.     [For  the  formation  of 
this  precipitate  dilute  the  NH4OH  with  10  vols.  of  H2O,  and 
use  only  one  or  two  drops  of  the  reagent.]     If  the  silver  solu- 
tion is  very  acid  no  precipitate  will  be  formed. 

3.  Sodium  Carbonate  precipitates  light  yellow  Ag2CO3.    Insol- 
uble in  excess.     Soluble  in  NH4OH  and  in  (NH4)2CO3. 


4  QUALITATIVE   ANALYSIS 

4.  Hydrogen    or    Ammonium    Sulfid    precipitates    black    Ag2S. 
Soluble  in  HNO3.     Insoluble  in  NH4OH. 

5.  Acid  Sodium  Phosphate  precipitates  yellow  Ag3PO4.     Sol- 
uble in  NH4OH  and  in  HNO3. 

6.  Potassium  Cyanid  precipitates  white  AgCN.      Soluble   in 
excess,  forming  AgCN(KCN).     From  this  solution  HNO3  pre- 
cipitates AgCN. 

7.  Potassium  lodid  precipitates  light  yellow  Agl.     Only  very 
slightly    soluble    in    NH4OH,    but    easily    soluble    in    KCN. 
Insoluble  in  HNO3. 

8.  Potassium  Chromate  precipitates  red-brown  Ag2CrO4.     Sol- 
uble in  HNO3  and  in  NH4OH. 

9.  Potassium   Sulfocyanate  precipitates  white  AgSCN.     Sol- 
uble in  NH4OH. 

10.  Potassium    Ferrocyanid    precipitates     white     Ag4Fe(CN)6. 
Difficultly  soluble  in  NH4OH. 

11.  Hydrochloric  Acid,  or  any  Soluble  Chlorid,  precipitates  white 
AgCl.   Soluble  in  NH4OH,  forming  (NH3)3(AgCl)2,  and  in  KCN. 
From  these  solutions  HNO3  reprecipitates  the  AgCl. 

12.  Metallic  Zinc,  Copper,  or  Mercury  will  precipitate  the  silver 
in  crystalline  form. 

13.  Reducing  Agents,  such  as  sulfurous  acid,  stannous  chlorid, 
or  ferrous  sulfate,   will  precipitate  the  silver  as  a  fine  gray 
powder. 

MERCURY  (Mercurous),  Hg' 

Mercury  dissolves  easily  in  HNO3.  If  the  mercury  is  in 
excess  there  is  formed  mercurous  nitrate,  HgNO3;  but  if  the 
HNO3  is  in  excess,  mercuric  nitrate,  Hg(NO3)2,  is  formed. 

Mercury  dissolves  in  hot  concentrated  H2SO4,  forming  mer- 
curic sulfate,  HgSO4,  and  SO2.  It  is  insoluble  in  HC1. 

For  the  reactions  use  a  solution  of  HgNO3. 


REACTIONS  FOR  THE  METALS  IN   SOLUTION  5 

1.  Sodium  Hydroxid  precipitates  black  Hg2O.    Insoluble  in  ex- 
cess of  the  reagent.    Decomposed  by  boiling  into  HgO  and  Hg. 

2.  Ammonium   Hydroxid,   or   Ammonium   Carbonate,  precipitates 
a  black  mixture  of  amido-mercuric  nitrate,  HgNH2NO3,  and 
finely  divided  mercury. 

3.  Sodium     Carbonate    precipitates    a    brownish-black    basic 
carbonate. 

4.  Hydrogen  or  Ammonium  Sulfid  precipitates  black  HgS  mixed 
with  Hg.     Soluble  in  aqua  regia.     If  this  precipitate  is  boiled 
with    concentrated  HNO3  it    is   changed   into    a   white   basic 
compound,  Hg(NO3)2(HgS)2. 

5.  Potassium  lodid  precipitates  yellowish-green  Hgl.  If  an 
excess    of    the    reagent    is    added    potassium    mercuric  iodid, 
HgI2(KI)2,    is    formecl.     This    dissolves    in    the    liquid,  while 
metallic  mercury  is  precipitated  as  a  gray  powder. 

6.  Potassium  Chromate  precipitates  brick-red  Hg2CrO4.    Diffi- 
cultly soluble  in  HNO3. 

7.  Hydrochloric   Acid,   or  a    Soluble  Chlorid,  precipitates  white 
HgCl  (calomel).     The  addition  of  NH4OH  changes  this  precipi- 
tate to  a  black  mixture  of  amido-mercuric  chlorid,  HgNH201, 
and  finely  divided  mercury. 

8.  Sulfuric  Acid,  in  not  too  dilute  solutions,  precipitates  white 
Hg2S04. 

9.  Stannous  Chlorid,  in  very  small  quantity,  precipitates  white 
HgCl.     In  excess  of  the  reagent  the  precipitate  is  reduced  to 
gray  metallic  mercury. 

10.  Metallic  Copper,  or  Zinc,  in  solutions  slightly  acidified  with 
HC1,  precipitates  metallic  mercury. 

11.  Sulfurous  Acid  reduces  a  mercurous  solution  to  metallic 
mercury,  which  can  often  be  collected  in  a  globule  by  boiling 
with  HCL 


6  QUALITATIVE   ANALYSIS 

MERCURY  (Mercuric),  Hg" 

Mercury  and  most  mercurous  compounds  can  be  changed  to 
mercuric  by  heating  with  concentrated  HNO3. 

For  the  reactions  use  a  solution  of  HgCl2,  or  Hg(NO3)2. 

1.  Sodium  Hydroxid  precipitates  yellow  Hg(OH)2. 
Insoluble  in  excess.     Soluble  in  warm  acids.     If  only  a  few 

drops   of   the   reagent  are   used    a  brown  basic   compound  is 
formed. 

2.  Ammonium  Hydroxid,  or  Ammonium  Carbonate,  precipitates  white 
amido-mercuric  chlorid,  HgNH2Cl. 

3.  Sodium    Carbonate    precipitates    a   brown    basic    carbonate, 
HgC03(HgO)8. 

4.  Hydrogen  or  Ammonium  Sulfid  precipitates  at  first  a  white 
double  salt,  HgCl2(HgS)2.     Excess  of  the  reagent  causes  this 
precipitate  to  change  to   yellow,    orange,  and  finally  to  black 
HgS.     Insoluble  in  concentrated  HNO3.     Long-continued  boil- 
ing with  HNO3  changes  it  into  a  white  basic  compound.     [See 
Mercurous  4.] 

5.  Potassium  lodid  precipitates  scarlet-red  HgI2.     Easily  sol- 
uble in  excess,   forming   HgI2(KI)2.     The   precipitate   is  first 
yellow,  then  salmon  red,  and  finally  scarlet  red. 

6.  Potassium  Chromate  precipitates,  from  not  too  dilute  solu- 
tions, orange  red  HgCrO4. 

7.  Metallic  Copper  precipitates,  from  solutions  acidified  with 
HC1,  gray  metallic   mercury. 

8.  Stannous  Chlorid  precipitates  first  white  mercurous  chlorid, 
HgCl,  and  in  excess,  gray  metallic  mercury. 

Other  reducing  agents  produce  a  similar  change. 


REACTIONS  FOR  THE  METALS  IN  SOLUTION  7 

BISMUTH,  Bi'" 

Bismuth  dissolves  easily  in  HNO3,  in  hot  concentrated  H2SO4, 
and  in  aqua  regia,  but  not  in  HC1. 

The  ordinary  bismuth  salts  are  not  soluble  in  H2O  except  in 
the  presence  of  considerable  free  acid,  usually  HNO3  or  HC1. 

For  the  reactions  use  a  solution  of  Bi(NO3)3,  in  dilute  HNO3. 

1.  Water  in  large  quantities  precipitates,  if  too  much  free 
acid  is  not  present,  white  basic  bismuth  nitrate,  Bi(OH)2NO3. 

If  much  free  HNO3  is  present  the  addition  of  ammonium 
chlorid,  or  of  HC1,  will  cause  the  precipitation  of  white 
BiOCl.  These  precipitates  are  insoluble  in  tartaric  acid.  [See 
Antimony  1.] 

2.  Sodium  or  Ammonium  Hydroxid  precipitates  white  Bi(OH)3. 
Insoluble  in  excess.     Changed  by  boiling  to  yellow  Bi2O3. 

3.  Sodium  or  Ammonium  Carbonate  precipitates  white  basic  bis- 
muth carbonate  (BiO)2CO3. 

4.  Hydrogen  or  Ammonium  Sulfid  precipitates  dark  brown  Bi2S3. 
Insoluble  in  (NH4)2S.     Soluble  in  HNO3. 

5.  Acid  Sodium  Phosphate  precipitates  white  BiPO4.     Insoluble 
in  dilute  acids. 

6.  Potassium  lodid  precipitates  brown  BiI3.    Soluble  in  excess. 

7.  Potassium  Chromate  precipitates  yellow  basic  bismuth  chro- 
mate  2  [(BiO)2CrO4]Bi2O3.    Insoluble  in  NaOH.     [See  Lead  8.] 
Soluble  in  HNO3. 

8.  Sodium  Stannite,  which  is  formed  by  adding  NaOH  to  a 
solution  of  SnCl2  until  the  precipitate  first  formed  is  dissolved, 
reduces   the  bismuth  solution   and  forms  a  black  precipitate, 
which  is  a  mixture  of  Bi  and  Bi2O3 ;  or,  if  the  reagent  is  added 
in  large  excess,  and  hot,  it  precipitates  metallic  bismuth. 


8  QUALITATIVE  ANALYSIS 

COPPER,  Cu" 

Copper  dissolves  easily  in  HNO3,  forming  Cu(NO3)2  and  NO, 
and  in  hot  concentrated  H2SO4,  forming  CuSO4  and  S^2*  ^  is 
only  very  slightly  soluble  in  dilute  H2SO4  or  HC1. 

For  the  reactions  use  a  solution  of  CuSO4. 

1.  Sodium  Hydroxid  precipitates  light  blue  Cu(OH)2.      Insol- 
uble in  excess  of  the  reagent.     Soluble  in  NH4OH. 

If  the  precipitate  is  boiled  with  an  excess  of  the  reagent  it 
becomes  black,  owing  to  the  formation  of  Cu(OH)2(CuO)2. 

2.  Ammonium  Hydroxid  precipitates  a  light  blue  basic  salt.    Easily 
soluble  in  excess  of  the  reagent,  forming  CuSO4(NH3)4H2O, 
which  gives  a  deep-blue  color  to  the  solution  (a  very  character- 
istic reaction).     If  KCN  is  added  to  this  blue  solution  the  color 
disappears,  owing  to  the  formation  of  Cu(CN)2(KCN)2. 

3.  Sodium    Carbonate    precipitates    a    blue    basic    carbonate, 
Cu2(OH)2CO3.      On    boiling,   this    precipitate    loses    CO2   and 
forms  black  Cu(OH)2(CuO)2. 

4.  Hydrogen  or  Ammonium  Sulfid  precipitates  black  CuS.     Sol- 
uble in  KCN  and  in  HNO3.     Insoluble  in  dilute  H2SO4.     [See 
Cadmium  4.] 

5.  Acid  Sodium  Phosphate  precipitates  greenish-blue  Cu3(PO4)2. 
Soluble  in  NH4OH. 

6.  Potassium    Cyanid    precipitates    greenish-yellow    Cu(CN)9. 
Easily  soluble  in  excess  of  the  reagent,  forming  Cu(CN)2(KCN)2. 
From  this  solution  H2S  will  not  precipitate  the  copper.     [See 
Cadmium  6.] 

7.  Potassium   lodid   precipitates   white   Cu2T2   and   free   iodin. 
The  latter  colors  the  precipitate  brown.     If  H2SO3  is  added 
the  precipitate  appears  white. 


REACTIONS   FOR  THE  METALS  IN   SOLUTION  9 

8.  Potassium   Sulfocyanate   precipitates   black  Cu(SCN)2.     If 
H2SO3   is   added   in   excess   the   copper  is  reduced  and  white 
Cua(SCN)2  is  formed. 

9.  Po.  jsium  Ferrocyanid  precipitates  red-brown  Cu2Fe(CN)6. 
[This  is  an  exceedingly  delicate  reaction,  one  part  of  copper 
showing  a  reddish  coloration  in  200,000  parts  of  water.] 

10.    Metallic  Iron  precipitates  copper  from  a  solution. 


CADMIUM,  Cd" 

Cadmium  dissolves   easily  in  HNO3,  and  slowly  in  H2SO4 
and  HC1. 

For  the  reactions  use  a  solution  of  Cd(NO3)2. 

1.  Sodium  Hydroxid  precipitates  white  Cd(OH)2.     Insoluble 
in  excess  of  the  reagent. 

2.  Ammonium    Hydroxid    precipitates    the    same    compound. 
Easily  soluble  in  excess  of  the  reagent. 

3.  Sodium  or  Ammonium  Carbonate  precipitates  white  CdCO3. 
Insoluble  in  excess  of  the  reagent.      Soluble  in  NH4OH. 

4.  Hydrogen    or  Ammonium    Sulfid    precipitates    yellow   CdS. 
Insoluble   in   (NH4)2S   or  in   KCN.      Soluble  in    warm  dilute 
H2SO4  or  HNO8.      [See  Copper  4.] 

5.  Acid  Sodium  Phosphate  precipitates  white  Cd3(PO4)2.     Sol- 
uble in  NH4OH  and  in  dilute  acids. 

6.  Potassium  Cyanid  precipitates  white  Cd(CN)2.     Soluble  in 
excess  of  the  reagent,  with  formation  of  Cd(CN)2(KCN)2.    From 
this  solution  H2S  precipitates  CdS.     [See  Copper  6.] 

7.  Potassium  Chromate  precipitates  a  yellow  basic  chromate 
Cd2(OH)2CrO4.      Since   this   precipitate    forms    slowly,   use  a 
slight  excess  of  the  reagent  and  allow  it  to  stand  for  a  few 
minutes.      Insoluble  in  NaOH.      [See  Lead  8.] 


10  QUALITATIVE   ANALYSIS 


ARSENIC,  As'" 

This  element  exists  in  both  trivalent  and  pentavalent  rela- 
tions, and  has  very  few  metallic  properties.  It  does  not 
dissolve  in  the  acids  to  form  salts,  and  we  have  already  learned 
that  it  forms  acids  quite  analogous  to  those  of  phosphorus. 
The  principal  reactions  of  arsenic  are  therefore  to  be  found 
among  those  of  the  acids. 

Arsenious  oxid,  As2O3,  dissolves  in  HC1,  and  this  solution 
may  be  used  for  the  reactions. 

1.  Hydrogen  Sulfid  precipitates  yellow  As2S3.  Insoluble  in  HC1. 
Soluble  in  NH4OH  or  (NH4)2CO3.  It  dissolves  in  (NH4)2S, 
forming  ammonium  sulfarsenite,  (NH4)3AsS3,  and  in  yellow 
ammonium  sulfid,  (NH4)2SX,  forming  ammonium  sulfarsenate, 
(NH4)3AsS4.  HC1  precipitates  from  these  solutions,  in  the 
first  case  As2S3,  in  the  second  As2S5. 

The  other  reagents  for  the  metals  give  no  precipitates  with 
arsenic  solutions. 

If  a  solution  containing  pentavalent  arsenic  is  treated  with 
H2S  it  is  reduced,  and  the  arsenic  precipitated  as  As2S3,  together 
with  sulfur.  This  action  takes  place  very  slowly  in  a  cold 
solution,  but  is  immediate  if  the  solution  is  hot. 

A  solution  of  sodium  arsenate,  Na3AsO4,  may  be  used  for 
this  reaction. 

ANTIMONY  (Stibium),  Sb'" 

Antimony  forms  both  trivalent  and  pentavalent  compounds. 
It  does  not  dissolve  in  HC1.  With  hot  concentrated  H2SO4  it 
forms  Sb2(SO4)3.  With  dilute  HNO3  it  forms  Sb2O3,  and  with 
concentrated  HNO3  it  forms  metantimonic  acid,  HSbO3.  It 
dissolves  in  aqua  regia,  forming  SbCl3  or  SbCl5,  according  to 
the  degree  of  concentration  of  the  acids  and  the  duration  of 
the  action. 


REACTIONS  FOR  THE  METALS  IN  SOLUTION  11 

These  compounds  do  not  dissolve  in  water  unless  free  hydro- 
chloric or  tartaric  acid  is  present. 

For  the  reactions  use  a  solution  of  SbCl3. 

1.  Water,    added    in    excess,    precipitates    white    antimony 
oxychlorid,  SbOCl.     Soluble  in  tartaric  acid,  so  that  if  much 
of  this  acid  is  present  the  precipitation  may  not  take  place. 

2.  Sodium  Hydroxid  precipitates  white   SbOOH.      Soluble  in 
excess  of  the  reagent,  with  formation  of  sodium  metantimonite, 
NaSb02. 

3.  Ammonium  Hydroxid,  or  Sodium  or  Ammonium  Carbonate,  pre- 
cipitates the  same.     Insoluble  in  excess  of  the  reagent. 

4.  Hydrogen  Sulfid  precipitates  orange-red   Sb2S3.      Insoluble 
in  (NH4)2CO3  [See  Arsenic  1]  and  in  dilute  acids.     Soluble  in 
warm   concentrated   II Cl.      Soluble   also   in   (NH4)aS,   forming 
ammonium  sulfantimonite,  (NH4)3SbS3,  and  in  (NH4)2SX,  form- 
ing ammonium  sulfantimonate,  (NH4)3SbS4.     HC1  precipitates 
from  these  solutions  Sb2S3  and  Sb2S5  respectively. 

5.  Metallic  Zinc,  in  solutions  containing  free  HC1,  precipitates 
the  antimony  as  a  black  powder.     If  a  piece  of  platinum  foil 
is  placed  in  the  solution,  in  contact  with  the  zinc,  the  antimony 
will  be  precipitated  on  the  foil  as  a  black  stain. 

Antimony,  in  its  pentavalent  relations,  is  acid  in  its  prop- 
erties. From  such  compounds  the  antimony  may  be  precipi- 
tated by  hydrogen  sulfid  as  orange-red  Sb2S5.  This  dissolves 
in  warm  concentrated  HC1,  forming  SbCl3  and  precipitating 
sulfur. 

TIN  (Stannous),  Sn" 

Tin  forms  both  stannous  (Sn")  and  stannic  (Sn"")  com- 
pounds. It  dissolves  in  HC1,  forming  SnCl2 ;  and  in  H2SO4, 
forming  SnSO4.  With  very  dilute  HNO3  it  forms  Sn(NO3)2, 
some  of  the  acid  being  reduced  forming  NH4NO3.  Thus : 

4  Sn  +  10  HN03  =  4  Sn(N03)2  +  NH4N03  +  3  H20. 


12  QUALITATIVE   ANALYSIS 

With  concentrated  HNO3  it  forms  white  stannic  acid,  H2SnO3. 
It  dissolves  in  aqua  regia,  forming  SnCl4. 
For  the  reactions  use  a  solution  of  SnCl2. 

1.  Sodium  Hydroxid  precipitates  white  Sn(OH)2.     Soluble  in 
excess  of  the  reagent,  forming  sodium  stannite,  Na2SnO2. 

2.  Ammonium  Hydroxid,  or  Sodium  or  Ammonium  Carbonate,  pre- 
cipitates the  same.      Insoluble  in  excess  of  the  reagent. 

3.  Hydrogen  or  Ammonium  Sulfid  precipitates  dark  brown  SnS. 
Insoluble  in  (NH4)2CO3  [See  Arsenic   1],   and  in  (NH4)2S   if 
free    from    (NH4)2SX.       Soluble    in    NaOH,    in    HC1,    and   in 
(NH4)2SX,  forming  with    the   latter    ammonium  sulfostannate, 
(NH4)2SnS3.     From  this  solution  HC1  precipitates  yellow  stan- 
nic sulfid,  SnS2. 

Stannous  chlorid,  and  all  other  stannous  compounds,  are 
easily  oxidized  to  stannic  compounds.  They  act,  therefore,  as 
powerful  reducing  agents  when  in  the  presence  of  reducible 
compounds.  Silver  salts  are  reduced  to  metallic  silver,  and  mer- 
cury salts  to  metallic  mercury.  [See  Silver  13,  Mercurous  9, 
and  Mercuric  8.]  Bismuth  compounds  are  reduced  to  metallic 
bismuth.  [See  Bismuth  8.]  Ferric  compounds  are  changed 
to  ferrous  compounds.  [See  Ferric  10.]  Potassium  chromate, 
K2CrO4,  and  potassium  permanganate,  KMnO4,  are  reduced  to 
chromium  chlorid  and  manganese  chlorid  respectively. 

2  K2Cr04  +  16  HC1  +  3  SnCl2  = 

4  KC1  +  2  CrCls  +  3  SnCl4  +  8  H20. 

Many  other  compounds  give  a  similar  reaction. 

TIN  (Stannic),  Sn"" 

Stannic  compounds  decompose  on  standing  and  precipitate 
Sn(OH)4.  If  stannous  chlorid  be  acidified  with  HC1,  a  few 
crystals  of  potassium  chlorate  added,  and  the  whole  boiled  until 
the  chlorous  odors  are  driven  away,  the  SnCl4  thus  formed  may 
be  used  for  the  following  reactions. 


REACTIONS   FOR  THE  METALS  IN  SOLUTION  13 

1.  Sodium  Hydroxid  precipitates  white  stannic  acid,  H2SnO3. 
Soluble   in  excess   of  the  reagent,   forming  sodium   stannate, 
Na2SnO3.     Soluble  also  in  the  mineral  acids. 

2.  Ammonium  Hydroxid,  or  Sodium  or  Ammonium  Carbonate,  pre- 
cipitates the  same.     Insoluble  in  excess  of  the  reagent. 

3.  Hydrogen  Sulfid  precipitates  from   solutions  which  do  not 
contain  too  large  an  excess  of  HC1,  yellow  SnS2.     Soluble  in 
concentrated  HC1,  and  in  (NH4)2S,  forming  ammonium  sulfostan- 
nate,  (NH4)3SnS3.     Insoluble  in  (NH4)2CO3.     [See  Arsenic  1.] 

4.  If  a  solution  of  stannic  chlorid  is  boiled  in  the  presence  of 
some  neutral  salt,  such  as  sodium  sulfate  or  ammonium  nitrate, 
metastannic  acid,  H10Sn5O15,  is  precipitated.     This  compound 
is  a  polymeric  form  of  stannic  acid,  and  is  insoluble  in  HNO3 
or  H2SO4. 

ALUMINUM,  Al'" 

Aluminum  dissolves  easily  in  HC1,  with  some  difficulty  in 
H2SO4,  and  scarcely  at  all  in  HNO3.  It  dissolves  also  in  NaOH 
and  in  KOH,  liberating  hydrogen. 

For  the  reactions  use  a  solution  of  A12(SO4)3. 

1.  Sodium  Hydroxid  precipitates  white  A1(OH)3.     Easily  soluble 
in  excess  of  the  reagent,  forming  sodium  aluminate,  NaAlO2. 
From  this  solution  it  is  reprecipitated  by  NH4C1.     Soluble  in 
all  mineral  acids  and  in  acetic  acid.* 

2.  Ammonium  Hydroxid  precipitates  the  same.     Very  slightly 
soluble  in  excess  of  the  reagent,  but  reprecipitated  by  boiling. 
If  NH4C1  is  present  the  precipitate  is  not  dissolved  in  excess  of 
the  reagent.* 

3.  Sodium  or  Ammonium  Carbonate  precipitates  the  same,  liber- 
ating CO2.* 

4.  Ammonium  Sulfid  precipitates  the  same,  liberating  H2S.* 

*  The  presence  of  non-volatile  organic  substances,  such  as  tartaric  acid,  citric 
acid,  sugar,  etc.,  prevents  this  precipitation. 


14  QUALITATIVE  ANALYSIS 

5.  Acid  Sodium  Phosphate  precipitates  white  A1PO4.     Soluble 
in  mineral  acids  and  in  NaOH.     From  the  solution  in  NaOH, 
NH4C1  precipitates  the  aluminum  as  A1(OH)3. 

6.  Sodium  Acetate  gives  no  precipitate  if  the  solution  is  cold, 
but  if  added  in  large  excess  and  boiled  the  aluminum  is  com- 
pletely precipitated  as  basic  aluminum  acetate,  A1(OH)2(C2H3O2). 

The    solution    must    be    neutral.     If    acid,    neutralize    with 
Na2CO3  or  NaOH. 

7.  Barium  Carbonate  precipitates  white  A1(OH)3,  liberating  CO2. 

CHROMIUM,  Cr"' 

Chromium  dissolves  in  HC1  and  H2SO4,  but  is  not  soluble  in 
HN03. 

For  the  reactions  use  a  solution  of  Cr2(SO4)3. 

1.  Sodium  Hydroxid  precipitates  gray-green  Cr(OH)3.     Soluble 
in  excess  of  the  reagent,  giving  a  dark  green  solution  and  form- 
ing sodium  chromite,  NaCrO2.     Reprecipitated  by  boiling  or  by 
the  addition  of  NH4C1. 

2.  Ammonium  Hydroxid  precipitates  the  same.  ^The  precipitate 
is  slightly  soluble  in  excess  of  the  reagent  whe%  cold  and  con- 
centrated, giving  a  reddish  color  to  the  solution.     Reprecipi- 
tated by  boiling  or  by  the  addition  of  NH4C1. 

3.  Sodium  or  Ammonium  Carbonate  precipitates  the  same,  libera- 
ting CO2.     The  precipitate  often  contains  some  basic  chromium 
carbonate  of  variable  composition. 

4.  Ammonium  Sulfid  precipitates  the  same,  liberating  H2S. 

5.  Acid  Sodium  Phosphate  precipitates  gray-green  CrPO4.     Sol- 
uble in  the  mineral  acids  and  in  NaOH. 

6.  Sodium  Acetate  gives  no  precipitate  unless  iron  or  aluminum 
salts  are  present,  in  which  case  the  chromium  is  partially  pre- 
cipitated by  boiling. 


REACTIONS  FOR  THE  METALS   IN   SOLUTION  15 

7.  Barium  Carbonate  precipitates  gray-green  Cr(OH)3,  liberating 
CO2. 

All  chromium  compounds,  when  treated  with  suitable  oxi- 
dizing agents,  are  converted  into  compounds  of  chromic  acid. 
A  common  method  of  oxidation  is  to  heat  the  compound  on  a 
piece  of  platinum  foil  with  a  mixture  of  Na2CO3  and  KNO3, 
which  gives  the  following  result. 

2  Cr(OH)3  +  3  KN03  +  2  Na2C03  = 

2  Na2O04  +  3  KN02  +  3  H20  +  2  C02. 

The  reactions  for  chromic  acid  will  be  given  with  those  of 
the  other  acids. 

IRON  (Ferrous),  Fe" 

Iron  forms  both  ferrous  (Fe")  and  ferric  (Fe'")  compounds. 
It  dissolves  easily  in  HC1  or  in  H2SO4,  forming  FeCl2  and 
FeSO4  respectively.  It  dissolves  in  HNO3,  forming  Fe(NO3)3, 
and  in  aqua  regia,  forming  FeCl3. 

For  the  reactions  use  a  solution  of  FeCl2  or  FeSO4. 

1.  Sodium  Hydroxid  precipitates  white  Fe(OH)2.     Insoluble  in 
excess  of  the  reagent.     The  white  color  of  the  precipitate  may 
be  seen  in  a  freshly  reduced  solution,  but  only  for  a  moment, 
since  it  absorbs  oxygen  from  the  air,  changes  first  to  a  dirty 
green,  and  then  to  a  red-brown  color,  forming  Fe(OH)3.     If 
ammonium  salts  are  present  the  precipitation  is  not  complete. 

2.  Ammonium    Hydroxid    partially    precipitates    the    iron    as 
Fe(OH)2.     If  ammonium  salts  are  present  no  precipitate  appears 
at  first ;  but  on  standing  the  iron  is  precipitated  as  red-brown 
Fe(OH)3. 

3.  Sodium  or  Ammonium  Carbonate  precipitates,  under  the  same 
conditions  as  above,  white  FeCO3.     This  loses  CO2,  oxidizes 
very  easily,  and  is  slowly  changed  to  Fe(OH)3. 

4.  Hydrogen  Sulfid  gives  no  precipitate  in  acidified  solutions. 
In  a  neutral  solution  it  gives  a  partial  precipitation  of  the  iron 


16  QUALITATIVE  ANALYSIS 

as    black    FeS.     If  the  solution  contains  sodium  acetate   the 
precipitation  is  nearly  complete. 

5.  Ammonium  Sulfid  precipitates  black  FeS.    Easily  soluble  in 
the  mineral  acids.     Difficultly  soluble  in  acetic  acid.     The  pre- 
cipitate oxidizes  easily  when  exposed  to  the  air,  forming  FeSO4 
and  a  basic  ferric  sulfate. 

6.  Potassium  Cyanid  precipitates  light  brown  Fe(CN)2.   Soluble 
in  excess  of  the  reagent,  forming  Fe(CN)2(KCN)4  or  K4Fe(CN)6. 

7.  Potassium   Sulfocyanate  gives   no   coloration   unless   ferric 
salts  are  present.     [See  Ferric  6.] 

8.  Potassium  Ferrocyanid  precipitates  bluish-white  potassium 
ferrous  ferrocyanid,  K2Fe"3[Fe"(CN)6]2.     This  absorbs  oxygen 
from  the  air  and  quickly  becomes  blue.     [See  Ferric  7.] 

9.  Potassium     Ferricyanid     precipitates     "  Turnbull's    blue," 
Fe"3[Fe'"(CN)6]2.     Insoluble  in  HC1.     Decomposed  by  NaOH, 
forming  Fe(OH)2,  which  oxidizes  very  rapidly,  giving  Fe(OH)3. 

10.  Barium  Carbonate  does  not  precipitate  iron  from  ferrous 
solutions.  [See  Ferric  9.] 

IRON  (Ferric),  Fe"' 

When  iron  is  dissolved  in  HNO3,  or  in  aqua  regia,  or  when 
ferrous  salts  are  acted  upon  by  oxidizing  agents,  such  as  HNO3, 
or  KC1O3  and  HC1,  ferric  compounds  are  formed. 

6  FeCl2  4-  KC103  +  6  HC1  =  KC1  +  3  H20  +  6  FeCl3. 
For  the  reactions  use  a  solution  of  FeCl3. 

1.  Sodium    or    Ammonium    Hydroxid    precipitates     red-brown 
Fa(OH)3.     Insoluble  in  excess  of  the  reagent.     Soluble  in  any 
mineral  acid.     [If  any  non-volatile  organic  substance,  such  as 
tartaric  acid,  is  present,  NH4OH  gives  no  precipitate.] 

2.  Sodium  or  Ammonium  Carbonate  precipitates  the  same,  libera- 
ting CO2.     The  precipitation  is  only  complete  after  boiling. 


REACTIONS  FOR  THE  METALS  IN  SOLUTION  17 

3.  Hydrogen  Sulfid  reduces  ferric  salts  to  ferrous  salts,  giving 
a  lightrcolored  precipitate  of  sulfur. 

4.  Ammonium  Sulfid  reduces  ferric  salts  to  ferrous  salts,  and 
precipitates  black  FeS  and  sulfur.     The  FeS  is  soluble  in  HC1, 
the  sulfur  remaining  undissolved. 

5.  Acid  Sodium  Phosphate  precipitates  yellowish- white  FePO4. 
Soluble  in  HC1.     Insoluble  in  acetic  acid. 

6.  Potassium   Sulfocyanate  produces  a  blood-red  coloration  in 
the  solution,  owing  to  the  formation  of  Fe(SCN)3.     [A  very  deli- 
cate reaction.]     This  action  does  not  take  place  in  the  presence  of 
sodium  acetate  unless  HC1  is  added  in  excess.     [See  Ferrous  7.] 

7.  Potassium  Ferrocyanid  precipitates  "  Prussian"  or  "  Berlin 
blue,"  Fe'"4[Fe"(CN)6]3.    [A  very  characteristic  reaction.]    Insol- 
uble in  the  mineral   acids.     Decomposed    by  NaOH,  forming 
red-brown  Fe(OH)3. 

8.  Sodium  Acetate  produces  a  red  coloration  caused  by  the 
formation  of  Fe(C2H3O2)3.      [If  mineral  acids  are  present  they 
must  be  neutralized.     This  can  be  done  best  with  Na2CO3.] 
On  diluting  this  solution  and  boiling,  the  iron  is  completely 
precipitated  as  red-brown  basic  ferric  acetate,  Fe(OH)2(C2H3O2). 

9.  Barium  Carbonate  precipitates  red-brown  Fe(OH)3,  libera- 
ting CO2.     [See  Ferrous  10.] 

10.  Stannous  Chlorid,  or  any  other  reducing  agent,  reduces 
ferric  salts  to  ferrous  salts. 

NICKEL,  Ni" 

Nickel  forms  nickelous  (Ni")  and  a  few  nickelic  (Ni'")  com- 
pounds. It  dissolves  slowly  in  HC1  and  in  H2SO4,  forming 
NiCl2  and  NiSO4  respectively,  and  readily  in  HNO3,  forming 
Ni(N03)2. 

For  the  reactions  use  a  solution  of  Ni(NO3)2. 


18  QUALITATIVE  ANALYSIS 

1.  Sodium  Hydroxid  precipitates  apple-green  Ni(OH)2.     Insol- 
uble in  excess  of  the  reagent.     Soluble  in  NH4C1.     If  sodium 
hypochlorite,  NaOCl,  or  bromin  water  with  excess  of  NaOH,  is 
added  to  this  precipitate,  it  is  oxidized  to  black  Ni(OH)3. 

2.  Ammonium  Hydroxid  precipitates  the  same  from  a  neutral 
solution.     Easily  soluble  in  excess  of  the  reagent  to  a  light 
blue  solution.     If  ammonium  salts  are  present,  or  some  free 
acid,  by  neutralizing  which  ammonium  salts  would  be  formed, 
no  precipitate  appears. 

3.  Sodium  or  Ammonium  Carbonate  precipitates  an  apple-green 
basic  carbonate  of  variable  composition.     Soluble  in  (NH4)2CO3 
to  a  blue  solution. 

4.  Hydrogen  Sulfid  gives  no  precipitate  if  the  solution  contains 
a  free  mineral  acid.     If  the  solution  is  neutral  it  gives  a  partial 
precipitation  of  black  NiS.     Sodium  acetate  added  in  excess  to 
the  nickel  solution  forms  nickel  acetate,  from  which  solution 
H2S  precipitates  all  the  nickel  as  black  NiS. 

5.  Ammonium  Sulfid  precipitates  the  same.     Slightly  soluble 
in  excess  of  the  reagent  to  a  dark  brown  solution,  from  which 
the  NiS  can  be  reprecipitated  by  boiling  or  by  the  addition  of 
acetic  acid.     Insoluble  in  dilute  HC1  or  acetic  acid.     Soluble 
in  warm  HNO3  or  in  aqua  regia. 

6.  Acid  Sodium  Phosphate  precipitates  apple-green  Ni3(PO4)2. 
Easily  soluble  in  dilute  acids. 

7.  Potassium  Cyanid  precipitates  yellow-green  Ni(CN)2.     Sol- 
uble in  excess  of  the  reagent,   forming  Ni(CN)2(KCN)2,  and 
reprecipitated  from  this  solution  by  dilute  HC1.     If  NaOH  is 
added  to  the  latter  solution,  and  then  bromin  water  in  excess, 
black   Ni(OH)3   is    precipitated,    liberating   cyanogen   bromid, 
CNBr.     (Poison!     Work  under  a  hood.)     [See  Cobalt  7.] 

8.  Potassium  Ferrocyanid  precipitates  green  Ni2Fe(CN)6.    Insol- 
uble in  dilute  acids. 


REACTIONS  FOR  THE  METALS  IN  SOLUTION  19 

9.    Potassium  Ferricyanid  precipitates  yellow-brown 
Ni3[Fe(CN)6]2.     Insoluble  in  dilute  acids. 

10.  Potassium  Nitrite  gives  no  precipitate  in  a  nickel  solution. 
[See  Cobalt  10.] 

COBALT,  Co" 

Cobalt  in  its  chemical  relations  very  closely  resembles  nickel. 
It  dissolves  in  the  mineral  acids,  forming  the  corresponding  salts. 
These  in  solution,  or  when  they  contain  water  of  crystallization, 
are  red,  but  on  losing  water  become  blue. 

For  the  reactions  use  a  solution  of  Co(NO3)2. 

1.  Sodium  Hydroxid  precipitates  a  blue  basic  salt.     If  excess 
of  the  reagent  is  added,  and  the  whole  boiled,  the  precipitate  is 
changed  to  red  Co(OH)2.     On  standing,  this  slowly  oxidizes 
to  brown  Co(OH)3.     Sodium  hypochlorite  and  brpmin  water 
give  reactions  similar  to  those  with  nickel. 

2.  Ammonium  Hydroxid  precipitates  from  a  neutral  solution  a 
blue  basic  salt.     Soluble  in  excess  of  the  reagent  to  a  red-brown 
solution.    If  ammonium  salts  are  present  no  precipitate  appears, 
but  the  solution  becomes  red  brown. 

3.  Sodium  or  Ammonium  Carbonate  precipitates  a  red-lilac  basic 
carbonate   of  variable   composition.     Soluble  in  excess  of  the 
(NH4)2CO3  to  a  red  solution,  which  slowly  becomes  brown  by 
oxidation.     If  acid  sodium  carbonate  is  used  as  the  reagent 
it  precipitates  normal  cobalt  carbonate,  CoCO3. 

4.  Hydrogen  Sulfid  gives  no  precipitate  if  the  solution  contains 
a  free  mineral  acid.     In  a  neutral  or  alkaline  solution,  or  in  one 
containing  sodium  acetate,  it  precipitates  black  CoS. 

5.  Ammonium   Sulfid  precipitates   black   CoS.     Insoluble   in 
excess    of    the    reagent,    in    dilute    HC1,  and   in    acetic    acid. 
Soluble  in  warm  HNO3  and  in  aqua  regia. 

6.  Acid  Sodium  Phosphate  precipitates  blue  Co3(PO4)2.     Sol- 
uble in  the  mineral  acids  and  in  NH4OH. 


20  QUALITATIVE  ANALYSIS 

7.  Potassium  Cyanid  precipitates  red-brown  Co(CN)2.     Soluble 
in  excess  of  the  reagent,  forming  Co(CN)2(KCN)4,  and  reprecipi- 
tated  from  this  solution  by  dilute  HC1.    If  the  solution  in  KCN 
is  boiled  for  some  time  potassium  cobalticyanid,  K3Co(CN)6,  is 
formed.     [Analogous    to    potassium    ferricyanid.]     HC1   gives 
no  precipitate  in  this  solution. 

If  NaOH  and  bromin  water  are  added  to  the  solution  in 
KCN  the  same  compound,  K3Co(CN)6,  is  formed,  and  cobalt 
is  not  precipitated.  [See  Nickel  7.]  (Since  the  commercial 
cobalt  salts  often  contain  traces  of  nickel,  a  very  slight  precipi- 
tate will  generally  be  formed.) 

8.  Potassium  Ferrocyanid  precipitates  bluish-green  Co2Fe(CN)6. 
Soluble  in  concentrated  HC1  to  a  blue-green  solution. 

9.  Potassium  Ferricyanid  precipitates   brown   Co3[Fe(CN)6]2. 
Insoluble  in  HC1. 

10.  Potassium  Nitrite,  added  in  excess  to  a  cobalt  solution 
which  has  been  previously  acidified  with  acetic  acid,  precipi- 
tates yellow  cobaltic-potassium  nitrite,  Co(NO2)3(KNO2)3.  The 
reaction  is  represented  by  the  following  equation : 

Co(N03)2  +  7  KN02  +  2  H(C2H302)  = 

Co(N02)3(KN02)3  4-  2  KN03  +  2  K(C2H302)  +  H2O  +  NO. 

The  precipitate  forms  slowly  in  dilute  solutions,  and  so  should 
be  allowed  to  stand  some  time.  The  precipitate  is  somewhat 
soluble  in  pure  water,  but  insoluble  in  the  presence  of  KNO2. 
[See  Nickel  10.] 

MANGANESE,  Mn" 

Manganese  forms  four  classes  of  compounds,  two  in  which  it 
is  basic,  and  two  in  which  it  is  acid.  These  are  the  manganous 
(Mn")  and  manganic  salts  (Mn'"),  manganates  (Mnvi)  and  per- 
manganates (Mnvii).  It  dissolves  easily  in  most  acids,  forming 
manganous  salts,  which  are  the  common  ones. 

For  the  reactions  use  a  solution  of  MnSO4. 


REACTIONS  FOR  THE  METALS  IN   SOLUTION  21 

1.  Sodium  Hydroxid  precipitates  white  Mn(OH)2.      Insoluble 
in  excess  of  the  reagent.     Soluble  in  NH4C1.     The  precipitate 
oxidizes  slowly  in  the  air,  forming  brown  Mn(OH)3. 

2.  Ammonium  Hydroxid  precipitates  the  same  in  a  neutral  solu- 
tion.    In  a  solution  containing  ammonium  salts,  or  a  free  acid, 
no  precipitate  is  formed  at  first ;  but  on  standing,  the  solution 
soon  oxidizes,  and  all  the  manganese  is  finally  precipitated  as 
brown  Mn(OH)3. 

3.  Sodium  or  Ammonium  Carbonate  precipitates  white  MnCO3. 
Boiling  makes  the  precipitation  complete. 

4.  Hydrogen  Sulfid  gives  no  precipitate  in  either  neutral  or 
acid  solutions. 

5.  Ammonium  Sulfid  precipitates  flesh-colored  MnS.     Soluble 
in  dilute  mineral  acids  and  in  acetic  acid.    Insoluble  in  NH4C1, 
in  the  presence  of  which  the  precipitation  is  complete. 

6.  Acid  Sodium  Phosphate  precipitates  white  Mn3(PO4)2.     Sol- 
uble in  the  mineral  acids  and  in  acetic  acid.     If  the  precipitate 
is  dissolved  in  HC1,  an  excess  of  NH4OH  added,  and  the  whole 
boiled,  a  light  rose-colored  crystalline  precipitate  of  MnNH4PO4 
is  formed. 

7.  Potassium     Ferrocyanid     precipitates     white     Mn2Fe(CN)6. 
Easily   soluble   in  H2SO4  and  in  HNO3,   and  with  difficulty 
in  HC1. 

8.  Potassium  Ferricyanid  precipitates  brown  Mn3[Fe(CN)6]2. 

9.  If  a  manganese  compound  is  heated  on  a  piece  of  platinum 
foil  with  a  mixture  of  Na2CO3  and  KNO3  it  will  be  oxidized, 
forming  green  sodium  manganate,  Na2MnO4.     Thus : 

MnS04  +  2  Na2C03  +  2  KN03  = 

Na2MnO4  +  Na2S04  +  2  KN02  +  2  CO2. 


22  QUALITATIVE  ANALYSIS 

If  the  green  mass  is  dissolved  in  water  with  the  addition  of  a 
few  drops  of  acetic  acid,  the  color  of  the  solution  will  change  to 
red,  owing  to  the  formation  of  sodium  permanganate,  NaMnO4, 
and  dark  brown  MnO2  will  be  precipitated. 

10.  If  a  small  quantity  of  red  lead,  Pb3O4,  is  placed  in  a 
test-tube  with  2  cc.  of  concentrated  HNO3,  a  few  drops  of  the 
manganese  solution  added,  and  the  whole  carefully  warmed,  the 
solution  becomes  red  from  the  formation  of  permanganic  acid. 

2  MnSO4  +  5  Pb304  +  26  HN03  = 

2  HMn04  +  2  PbS04  +  13  Pb(N03)2  +  12  H20. 

11.  Manganic    and   permanganic   acids  and  their  salts  are 
easily  reduced  to  manganous  salts  in  the  presence  of  reducing 
agents,  such  as  H2SO3,  H2S,  or  nascent  hydrogen. 

ZINC,  Zn" 

Zinc  dissolves  easily  in  most  acids,  forming  the  correspond- 
ing salt.  It  also  dissolves  in  NaOH,  forming  sodium  zincate, 
Na2ZnO2. 

For  the  reactions  use  a  solution  of  ZnSO4. 

1.  Sodium  Hydroxid  precipitates  white  Zn(OH)2.     Easily  sol- 
uble in  excess  of  the  re~agent,  forming  sodium  zincate,  Na2ZnO2. 

2.  Ammonium  Hydroxid  precipitates  the  same.     Easily  soluble 
in  excess  of  the  reagent,  forming  ZnSO4(NH3)4.     Ammonium 
salts  prevent  the  precipitation. 

3.  Sodium  Carbonate  precipitates  a  white  basic  carbonate  of 
variable  composition,  but  usually  Zn2(OH)2CO3.    Boiling  makes 
the  precipitation  complete.     If  acid  sodium  carbonate  is  used 
as  the  reagent  it  precipitates  normal  zinc  carbonate,  ZnCO3. 

4.  Ammonium  Carbonate  precipitates   the   same.     Soluble  in 
excess    of    the    reagent.     Ammonium   salts   prevent   the   pre- 
cipitation. 


REACTIONS   FOR  THE  METALS  IN  SOLUTION  23 

5.  Hydrogen  Sulfid  gives  no  precipitate  in  solutions  containing 
a  free  mineral  acid.     In  a  neutral  or  alkaline  solution,  or  one 
acidified  with  acetic  acid,  it  precipitates  white  ZnS. 

6.  Ammonium  Sulfid  precipitates  the  same  from  any  solution. 
If  NH4C1  is  present  the  precipitation  is  complete. 

7.  Acid  Sodium  Phosphate  precipitates  white  Zn3(PO4)2.     Sol- 
uble in  dilute  acids  and  in  NH4OH. 

8.  Potassium  Cyanid  precipitates  white  Zn(CN)2.     Soluble  in 
excess   of  the   reagent,   forming   Zn(CN)2(KCN)2.     From  this 
solution  (NH4)2S  precipitates  white  ZnS. 

9.  Potassium  Ferrocyanid  precipitates  white  Zn2Fe(CN)6.     In- 
soluble in  dilute  acids  and  in  NH4OH. 

10.  Potassium  Ferricyanid  precipitates  brownish-yellow 
Zn8[Fe(CN)6]2.  Soluble  in  HC1  and  in  NH4OH. 

MAGNESIUM,  Mg" 

Magnesium  dissolves  easily  in  all  acids,  forming  the  corre- 
sponding salts.  If  heated  in  the  air  it  takes  fire  quite  easily, 
and  burns  with  an  intensely  white  light,  forming  MgO. 

For  the  reactions  use  a  solution  of  MgSO4. 

1.  Sodium  Hydroxid  precipitates  white  Mg(OH)2.      Soluble 
in   NH4C1.     The   presence    of   ammonium  salts  prevents  the 
precipitation. 

2.  Ammonium  Hydroxid  gives  a  partial  precipitation  of  the 
same  in  a  neutral  solution.     If  an  ammonium  salt  or  a  free 
acid  is  present  no  precipitate  appears. 

3.  Sodium  Carbonate  precipitates  a  white  basic  carbonate  of 
variable  composition.      If  the  precipitate  is  boiled  it  has  the 
composition  Mg3(OH)2(CO3)2.     If  ammonium  salts  are  present 
no  precipitate  is  formed. 


24  QUALITATIVE    ANALYSIS 

4.  Ammonium   Carbonate   gives    no    precipitate    if    ammonium 
Halts  are  present. 

5.  Acid  Sodium  Phosphate  precipitates  white  MgHPO4,  which 
by  boiling  changes  to  Mg3(PO4)2.     If  NH4C1  is  added  to  the 
solution,  and  then  NH4OH  in  excess,  the  reagent  precipitates 
white  crystalline  MgNH4PO4.     This  precipitate  forms  slowly  in 
a  dilute  solution  and  is  complete  only  after  standing  some  hours. 

6.  Ammonium  Oxalate  gives  no  precipitate  in  dilute  solutions. 
In  concentrated  solutions  it  gives  a  white  precipitate  of  MgC2O4. 
Soluble  in  NH4C1.     [See  Calcium  6.] 

BARIUM,  Ba" 

The  metal  barium  has  little  practical  value  and  is  very 
difficult  to  obtain  in  the  metallic  state. 

It  oxidizes  easily  in  the  air  and  decomposes  water  at  the 
ordinary  temperature. 

Its  salts  are  easy  to  form  and  many  of  them  are  soluble  in 
water. 

For  the  reactions  use  a  solution  of  BaCl2. 

1.  Sodium  Hydroxid   precipitates,   if    the   solution   is   not  too 
dilute,  white  Ba(OH)2.     Somewhat  soluble  in  cold  water,  much 
more  so  in  hot. 

2.  Ammonium  Hydroxid  gives  no  precipitate  in  barium  solutions. 

3.  Sodium  or  Ammonium  Carbonate   precipitates  white  BaCO3. 
Somewhat  soluble  in  NH4C1.    Soluble  in  water  containing  CO2, 
forming  the  acid   carbonate,    BaH2(CO3)2.     The    precipitation 
can  be  made    complete   by  adding  NH4OH    in  slight  excess 
and  boiling. 

4.  Acid  Sodium  Phosphate  precipitates  a  white  acid  phosphate, 
BaHPO4.     If  NH4OH  is  present  it  forms  BaNH4PO4.     Sol- 
uble in  dilute  acids  and  reprecipitated  by  NH4OH. 


REACTIONS   FOR  THE  METALS   IN  SOLUTION  25 

5.  Potassium  Chromate  precipitates  yellow  BaCrO4.     Soluble 
in  HC1  and  HNO3.     Insoluble  in  NaOH  [See  Lead  8],  and  in 
acetic  acid  [See  Strontium  5], 

6.  Sulfuric   Acid,   or  any   soluble    sulfate,    precipitates   white 
BaSO4.      This    precipitation    takes    place    even    in    extremely 
dilute  solutions.      Insoluble   in  all  acids  and  alkalies.      [See 
Strontium  6.] 

7.  Ammonium  Oxalate  precipitates,  if  the  solution  is  not  too 
dilute,   white  BaC2O4.     Dilute   solutions   give  no   precipitate. 
[See  Calcium  6.]     Soluble  in  HC1  and  HNO3. 

8.  Hydrofluosilicic  Acid  precipitates  white  BaSiF6.     Somewhat 
soluble    in   water.     Insoluble   in  alcohol  and  in  dilute   acids. 
[See  Strontium  8.] 

9.  If  a  barium  compound  is  heated  on  a  platinum  wire  in 
an  oxidizing  flame  it  imparts  a  pale  green  color  to  the  flame. 
If  the  compound  is  a  chlorid,  or  is  moistened  with  HC1,  the 
color  is  more  distinct. 

STRONTIUM,  Sr" 

Strontium  very    closely  resembles    barium    in  its    chemical 
properties. 

For  the  reactions  use  a  solution  of  SrCl2. 

1.  Sodium  Hydroxid  precipitates   white   Sr(OH)2.     Somewhat 
soluble  in  water,  but  less  so  than  Ba(OH)2. 

2.  Ammonium  Hydroxid  gives  no  precipitate. 

3.  Sodium  or  Ammonium   Carbonate  precipitates   white   SrCO3. 
Its  properties  are  like  those  of  BaCO3.     [See  Barium  3.] 

4.  Acid  Sodium  Phosphate  precipitates  white   SrHPO4.      Like 
BaHPO4.     [See  Barium  4.] 

5.  Potassium  Chromate  gives  no  precipitate  at  first,  but  after  a 
time,  if  the  solution  is  neutral  and  not  too  dilute,  yellow  SrCrO4 


26  QUALITATIVE   ANALYSIS 

is  precipitated.  Insoluble  in  alcohol  even  when  dilute,  so  that 
if  alcohol  is  added  to  the  solution  the  precipitate  appears  at 
once.  Soluble  in  acetic  acid.  [See  Barium  5.] 

6.  Sulfuric  Acid,   or   any  soluble  sulfate,   precipitates    white 
SrSO4.     Slightly  soluble  in  water,  so  that  if  the  solution  is 
very    dilute    the    precipitate     does    not    appear    immediately. 
[See  Barium  6  and  Calcium  5.]     A  concentrated  solution  of 
Na2CO3  or  (NH4)2CO3  converts  it  into  SrCO3. 

7.  Ammonium  Oxalate  precipitates  white  SrC2O4.     Somewhat 
soluble  in  water,  but  less  so  than  BaC2O4. 

8.  Hydrofluosilicic  Acid  gives  no  precipitate  in  moderately  dilute 
solutions.     [See  Barium  8.] 

9.  If  a  strontium  compound  is  heated  on  a  platinum  wire  in 
an  oxidizing  flame  it  imparts   a  crimson   color  to  the  flame. 
If  the  compound  is  a  chlorid,  or  is  moistened  with  HC1,  the 
color  is  more  distinct. 

CALCIUM,  Ca" 

Calcium  very  closely  resembles  strontium  and  barium  in  its 
chemical  properties. 

For  the  reactions  use  a  solution  of  CaCl2. 

1.  Sodium  Hydroxid  precipitates  white  Ca(OH)2.     Slightly  sol- 
uble in  water,  but  much  less  so  than  Sr(OH)2  or  Ba(OH)2. 

2.  Ammonium  Hydroxid  gives  no  precipitate. 

3.  Sodium   or  Ammonium  Carbonate  precipitates  white   CaCO3. 
Its  properties  are  like  those  of  BaCO3.     [See  Barium  3.] 

4.  Acid   Sodium    Phosphate    precipitates    white    CaHPO4.      If 
NH4OH  is  present  the  normal  salt,  Ca3(PO4)2,  is  precipitated. 
Soluble  in  dilute  acids  and  reprecipitated  by  NH4OH. 

5.  Sulfuric   Acid,  or  any  soluble  sulfate,   precipitates   white 
CaSO4.     Somewhat  soluble  in  water,  so  that  if  the  solution 


REACTIONS  FOR  THE  METALS  IN  SOLUTION  27 

is  very  dilute  no  precipitate  appears.     A  concentrated  solution 
of  Na2CO3  or  (NH4)2CO3  converts  it  into  CaCO3. 

6.  Ammonium  Oxalate  precipitates,  even  from  very  dilute  solu- 
tions, white  CaC2O4.     Soluble  in  HC1  or  HNO3.     Insoluble  in 
water  and    acetic  acid.      [See   Barium   7.]      The   presence   of 
NH4OH  hastens  the   precipitation,   which,   if  the  solution  is 
cold  and  dilute,  is  complete  only  after  long  standing. 

7.  Hydrofluosilicic  Acid  gives  no  precipitate  even  if  alcohol  is 
added.     [See  Barium  8.] 

8.  If  a  calcium  compound  is  heated  on  a  platinum  wire  in  an 
oxidizing  flame  it  imparts  a  yellowish-red  color  to  the  flame. 
If  the  compound  is  a  chlorid,  or  is  moistened  with  HC1,  the 
color  is  more  distinct. 

THE  ALKALI   METALS 

Potassium  and  sodium  are  the  common  elements  belonging 
to  the  alkali  group  of  metals.  They  are  very  strong  bases,  and 
form  salts  with  every  acid  known.  The  salts  are  all  soluble  in 
water  to  some  extent,  and  so  form  no  precipitates  with  the  com- 
mon reagents,  most  of  which  are  compounds  of  these  metals. 
There  are  a  few  compounds  which  are  difficultly  soluble  in  water, 
and  these  are  precipitated  when  they  are  produced  in  sufficiently 
concentrated  solutions.  Most  of  the  salts  of  these  metals  are 
either  insoluble,  or  difficultly  soluble,  in  alcohol,  so  that  the 
addition  of  this  reagent  often  helps  the  formation  of  a  pre- 
cipitate. The  compound  radical  ammonium,  NH4,  forms  a  series 

of  compounds  analogous  to  those  of  potassium  and  sodium. 

• 

POTASSIUM  (Kalium),  K' 

For  the  reactions  use  a  solution  of  KC1. 

1.  Acid  Sodium  Tartrate,  HNa(C4H4O6),  precipitates,  if  the 
solution  is  not  too  dilute,  white  crystalline  HK(C4H4O6).  The 


28  QUALITATIVE   ANALYSIS 

precipitate  forms  slowly,  but  may  be  hastened  by  shaking.     [See 
Ammonium  2.] 

2.  Hydrofluosilicic  Acid  precipitates,  if  the  solution  is  not  too 
dilute,  white  K2SiF6.    Insoluble  in  dilute  acids  and  in  alcohol. 

3.  If  to  2  cc.  of  a  solution  of  sodium  nitrite  there  are  added 
1  cc.  of  acetic  acid  and  5  drops  of  a  solution  of  cobalt  nitrate 
a  deep  orange-yellow  liquid  is  formed.     This  precipitates  from 
the     potassium     solution     yellow     cobaltic-potassium     nitrite, 
Co(NO2)8(KNOa)8.     [See  Cobalt  10.] 

4.  Platinum  Chlorid  *  precipitates  from  neutral  or  slightly  acid 
solutions   yellow    K2PtCl6.     Soluble    in    100    parts    of    water. 
Insoluble  in  alcohol. 

All  potassium  salts,  when  heated  on  a  platinum  wire  in  an 
oxidizing  flame,  impart  a  reddish-violet  color  to  the  flame. 
This  color  appears  red  when  seen  through  a  blue  glass. 

SODIUM  (Natrium),  Na' 
For  the  reactions  use  a  solution  of  NaCl. 

1.  Hydrofluosilicic  Acid  precipitates,  after  long  standing  if  the 
solution  is  dilute,  or  upon  addition  of  alcohol,  white  Na2SiF6. 

2.  Acid  Potassium  Pyroantimonate,  K2H2Sb2O7,  precipitates,  in 
neutral  solutions   which    do   not  contain   other   metals,   white 
Na2H2Sb2O7.     The  precipitation  is  slow,  but  may  be  hastened 
by  shaking. 

3.  All  sodium  salts,  when  heated  on  a  platinum  wire  in  an 
oxidizing  flame,  impart  a  bright  yellow  color  to  the  flame.     This 
color  is  not  seen  through  blue  glass.     A  crystal  of  potassium 
bichromate  appears  colorless  in  this  yellow  light. 

*  The  reactions  with  platinum  chlorid  may  be  omitted. 


REACTIONS  FOR  THE  METALS  IN   SOLUTION  29 

AMMONIUM,   (NH4)' 

For  the  reactions  use  a  solution  of  NH4C1. 

1.  Sodium  Hydroxid,  or  any  soluble  base,  when  heated  with  an 
ammonium  compound,   decomposes  it,   liberating   NH3.     This 
may  be  recognized  by  its  characteristic  odor,  or  by  the  white 
clouds  of  NH4C1  which  are  formed   if  a  rod  moistened  with 
HC1  is  held  in  the  escaping  gas. 

2.  Acid  Sodium  Tartrate  precipitates,  if  the  solution  is  not  too 
dilute,  HNH4(C4H4O6).     The  precipitation  may  be  hastened  by 
shaking.      [See  Potassium  1.] 

3.  Platinum  Chlorid  precipitates,  from  neutral  or  slightly  acid 
solutions,  yellow  (NH4)2PtCl6.     Soluble  in  170  parts  of  water. 
Insoluble  in  alcohol. 

4.  Nessler's  Reagent,  a  solution  of  HgI2(KT)2  with  an  excess 
of  KOH,  precipitates,  even  from   extremely  dilute  solutions, 
brown  Hg2I(NH2)O.     This  reaction  is  best  shown  by  filling  a 
test-tube  nearly  full  of  water,  adding  two  or  three  drops  of  the 
ammonium  solution,  and  then  the  reagent. 


REACTIONS   FOR   THE   ACID   RADICALS   IN 
SOLUTION 

For  the  reactions  for  the  acid  radicals  it  is  better  to  use  solu- 
tions of  the  salts  derived  from  the  acids  rather  than  the  acids 
themselves,  although  the  latter  may  sometimes  be  used.  The 
salts  are  usually  neutral  in  their  action  on  litmus  paper,  while 
the  acids  and  most  acid  salts  turn  the  blue  litmus  paper  red. 
The  free  acids  may  be  further  distinguished  by  leaving  no 
residue  when  a  few  drops  are  evaporated  to  dryness  on  a  piece 
of  platinum  foil. 

The  reactions  are  similar  to  those  for  the  metals  except  that 
the  solutions  which  were  then  used  for  the  reactions  now 
become  the  reagents,  and  the  reagents  then  used  are  now  the 
solutions  for  the  reactions.  In  testing  for  the  acid  radicals, 
therefore,  only  a  few  of  the  more  characteristic  reactions  will 
be  given.  For  other  reactions  the  student  is  referred  to  those 
given  under  the  different  metals. 

The  student  must  always  consider  the  nature  not  only  of 
the  reagent  used  but  also  of  the  substance  in  the  solution.  If 
either  is  reducing  in  its  action  this  will  manifest  itself  in  the 
precipitation.  [See  Silver  13,  Mercurous  9,  Bismuth  8,  etc.] 
The  relation  which  the  metal  in  each  bears  to  the  acid  radical 
in  the  other  must  also  be  considered.  If  the  metal  in  the 
unknown  solution  forms  a  precipitate  with  the  acid  radical  of 
the  reagent  no  information  regarding  the  acid  radical  of  the 
unknown  substance  can  be  obtained  by  this  particular  reaction. 
This  does  not  often  occur,  but  when  it  does  the  other  reactions 
must  be  relied  upon  for  proving  the  constitution  of  the  acid 
radical. 

30 


REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION     31 

HYDROCHLORIC  ACID,  HC1 

For  the  reactions  use  a  solution  of  NaCl. 

1.  Lead  Acetate  precipitates  white  PbCl2.     Soluble  in  boiling 
water  or  in  a  large  quantity  of  cold  water,  so  that  if  the  solu- 
tion is  very  dilute  the  precipitate  may  fail  to  appear. 

2.  Silver  Nitrate  precipitates  white  AgCl.     Easily  soluble  in 
NH4OH  and  in  KCN,  and  reprecipitated  from  these  solutions 
by  HN03. 

3.  Mercurous  Nitrate  precipitates  white  HgCl.     By  the  addition 
of  NH4OH  this  precipitate  becomes  black,  owing  to  the  forma- 
tion of  amido-mercuric  chlorid,   HgNH2Cl,  mixed  with  finely 
divided  mercury. 

The  changing  of  this  precipitate  from  white  to  black  by  the 
addition  of  NH4OH  is  characteristic  of  mercurous  compounds 
rather  than  of  chlorids. 

HYDROBROMIC  ACID,  HBr 

For  the  reactions  use  a  solution  of  KBr. 

1.  Lead  Acetate  precipitates  white  PbBr2.     Somewhat  soluble 
in  water,  but  not  as  easily  soluble  as  PbCl2. 

2.  Silver  Nitrate  precipitates  yellowish-white  AgBr.     Soluble 
with  some  difficulty  in  NH4OH,  but  easily  soluble  in  KCN. 
Insoluble  in  dilute  acids. 

3.  Mercurous  Nitrate  precipitates  yellowish-white  HgBr.     The 
precipitate  becomes  black  on  adding  NH4OH. 

4.  Chlorin  Water  liberates  bromin  from  many  of  its  compounds, 
coloring  the  solution   red  brown.     If  a  little  carbon  disulfid, 
CS2,  is  added  to  the  solution,  and  the  whole  well  shaken,  the 
bromin  dissolves  in  the  CS2  and  colors  it  red  brown. 


32  QUALITATIVE   ANALYSIS 

HYDRIODIC  ACID,  HI 
For  the  reactions  use  a  solution  of  KI. 

1.  Lead  Acetate  precipitates  yellow  PbI2.     Soluble  in  boiling 
water,  from  which  solution  it  crystallizes,  on  cooling,  in  golden 
yellow  scales. 

2.  Silver   Nitrate   precipitates  light  yellow  Agl.     Only  very 
slightly  soluble  in  NH4OH,  but  easily  soluble  in  KCN.    Insoluble 
in  HNO3. 

3.  Mercurous  Nitrate  precipitates  yellowish-green   Hgl.      [See 
Mercurous  5.] 

4.  Mercuric  Chlorid  precipitates  scarlet-red  HgI2.     Soluble  in 
KI.     [See  Mercuric  5.] 

5.  Bismuth  Nitrate  precipitates  brown  BiI3. 

6.  Copper  Sulfate  precipitates  white  Cu2I2  together  with  free 
iodin,  which  colors  the  precipitate  brown.     If  H2SO3  is  added 
the  precipitate  appears  white. 

7.  Chlorin  or  Bromin  Water  liberates  iodin  from   most  of  its 
compounds.     If  a  little  CS2  is  added  to  the  solution  and  the 
whole  well  shaken,  the  iodin  dissolves  in  the  CS2  and  colors 
it  violet.      If  starch  paste  is  added  to  the  solution,  the  iodin 
colors  it  deep  blue. 

HYDROFLUORIC  ACID,  HF 
For  the  reactions  use  a  solution  of  KF. 

1.  Lead  Acetate  precipitates  white  PbF2.     Soluble  in  HNO3. 

2.  Silver  Nitrate  gives  no  precipitate.      (Distinction   between 
fluorids  and  the  other  halogen  salts.) 

3.  Barium  Chlorid  precipitates  white  BaF2.     Soluble  in  HC1 
or  in  HNOQ. 


REACTIONS  FOR  THE  ACID  RADICALS  IN   SOLUTION     33 

4.  Calcium  Chlorid  precipitates  white  CaF2.     Scarcely  soluble 
in  any  dilute  acid. 

5.  All  fluorids  are  decomposed  by  concentrated  H2SO4,  lib- 
erating HF.     This  acid  unites  with  the  silicon  in  glass,  forming 
SiF4.    Hence  if  a  fluorid,  together  with  some  concentrated  H2SO4, 
is  heated  for  a  moment  in  a  clean  test-tube,  and  the  tube  then 
emptied  and  cleaned,  it  will  be  found  to  have  been  etched.     To 
show  this  reaction  the  solution  must  be  fairly  concentrated. 

HYDROCYANIC  ACID,  HCN 

For  the  reactions  use  a  solution  of  KCN. 

1.  Lead    Acetate    precipitates    white    Pb(CN)2.     Insoluble    in 

KCN. 

2.  Silver  Nitrate  precipitates  white  AgCN.     Soluble  in  KCN 
and   in   NH4OH,   and  reprecipitated   from   these   solutions   by 
HNO3. 

3.  Copper  Sulfate  precipitates  greenish-yellow  Cu(CN)2.     Sol- 
uble in  KCN,   from  which  solution   H2S  will  not  precipitate 
the  copper. 

4.  Cadmium  Nitrate  precipitates  white  Cd(CN)2.     Soluble  in 
KCN,  from  which  solution  H2S  precipitates  yellow  CdS. 

5.  If  a  few  drops  of  (NH4)2SX  are  added  to  a  solution  of 
KCN,  and  the  solution  boiled  for  a  moment,  potassium  sulfocy- 
anate,  KSCN,  is  formed.     If  HC1  is  now  added  in  excess, ferric 
chlorid    will    produce    a   blood-red    coloration,    owing    to    the 
formation  of  Fe(SCN)3. 

6.  If  a  small  quantity  of  NaOH  is  added  to  a  solution  of 
KCN,   then   three   or   four  drops   each   of   FeSO4   and  FeCl3, 
and    finally   HC1   in   excess,   Prussian    blue   is   formed.     [See 
Ferric  7.] 


34  QUALITATIVE  ANALYSIS 

SULFOCYANIC  OR  THIOCYANIC  ACID,  HSCN 

For  the  reactions  use  a  solution  of  KSCN. 

1.  Silver    Nitrate    precipitates    white    AgSCN.     Soluble    in 
NH4OH. 

2.  Mercurous  Nitrate  produces  a  gray  precipitate,  which  is  a 
mixture  of  Hg  and  Hg(SCN)2,  with  perhaps  some  HgSCN. 

3.  Mercuric  Nitrate  precipitates  white  Hg(SCN)2. 

4.  Copper  Sulfate  precipitates,  from  a  concentrated  solution, 
black  Cu(SCN)2.     If  the  solution  is  dilute,  an  emerald-green 
coloration  is  produced.     If  H2SO3  in  excess  is  added  to  this 
solution,  and  the  whole  boiled,  the  copper  is  reduced  and  white 
Cu2(SCN)2  is  precipitated.     [See  Copper  8.] 

5.  Ferric  Chlorid  produces  a  blood-red  coloration  in  the  solu- 
tion,  owing  to  the   formation   of  Fe(SCN)3.     This  is  a  very 
characteristic  reaction. 

HYDROFERROCYANIC  ACID,  H4Fe(CN)6 
For  the  reactions  use  a  solution  of  K4Fe(CN)6. 

1.  Lead  Acetate  precipitates  white  Pb2Fe(CN)6. 

2.  Silver  Nitrate  precipitates  white  Ag4Fe(CN)6.     Insoluble  in 
dilute  NH4OH. 

3.  Copper  Sulfate  precipitates  red-brown  Cu2Fe(CN)6.      This 
reaction  can  be  shown  in  a  very  dilute  solution. 

4.  Ferric  Chlorid  precipitates   Prussian   blue,   Fe4[Fe(CN)6]3. 
Decomposed  by  NaOH,  forming  red-brown  Fe(OH)3. 

HYDROFERRICYANIC  ACID,  H3Fe(CN)6 
For  the  reactions  use  a  solution  of  K3Fe(CN)6. 

1.    Silver  Nitrate  precipitates  red-brown  Ag3Fe(CN)6.     Soluble 
in  NH4OH. 


REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION     35 

2.  Ferrous  Sulfate  precipitates  Turnbull's  blue,  Fe8[Fe(CN)6]2. 
Insoluble  in  HCL 

3.  Ferric  Chlorid  gives  a  red-brown  solution  but  no  precipitate. 

4.  Zinc   Sulfate   precipitates   brownish-yellow    Zn3[Fe(CN)6]2. 
Soluble  in  HC1  and  in  NH4OH. 

HYPOCHLOROUS  ACID,  HC10 

All  hypochlorites  are  soluble  in  water,  and  so  the  acid  radical 
cannot  be  precipitated.  If  a  concentrated  solution  of  a  hypo- 
chlorite  is  boiled,  oxygen  is  liberated.  If  a  dilute  acid  is  added 
to  the  solution,  chlorin  is  liberated. 

For  the  reactions  use  a  solution  of  NaClO. 

1.  Lead  Acetate,  to  which  NaOH   has  been  added  until  the 
precipitate  first  formed  is  dissolved,  precipitates  brown  PbO2. 

Pb(C2H302)2  +  2  NaOH  +  NaClO  = 

Pb02  +  2  NaC2H302  +  NaCl  +  H20. 
The  precipitation  is  hastened  by  boiling. 

2.  Silver  Nitrate   gives   a   white   precipitate   of   AgCl,   silver 
chlorate  being  formed  at  the  same  time. 

3  AgN03  +  3  NaClO  =  2  AgCl  +  AgC103  +  3  NaN03. 

3.  If  a  piece  of  litmus  paper  is  moistened  with  a  few  drops 
of  the  solution  and  then  exposed  to  acid  fumes,  the  color  will 
be  bleached.     If   the   moistened   paper  is   breathed  upon,   the 
CO2  in  the  breath  will  effect  the  same  change. 

CHLORIC  ACID,  HC103 

The  chlorates  are  all  soluble  in  water  and  so  form  no  pre- 
cipitates. 

For  the  reactions  use  a  solution  of  KC1O3. 

1.  Silver  Nitrate  gives  no  precipitate  with  a  chlorate,  but  if 
H2SO3  is  added  to  the  solution,  the  chlorate  is  reduced  to  a 
chlorid,  and  AgNO3  then  gives  a  white  precipitate  of  AgCl. 


36  QUALITATIVE   ANALYSIS 

2.  Hydrochloric  Acid  decomposes  the  chlorates,  giving  chlorin 
peroxid  and  chlorin.  Thus  : 

2  HC1  +  KC1O3  =  KC1  +  H20  +  C102  +  Cl. 

The  chlorin  peroxid  and  chlorin  dissolve  in  the  solution, 
coloring  it  yellow;  but  if  the  solution  is  boiled  these  gases  will 
pass  off,  giving  what  is  called  a  "  chlorous  odor."  This  mixture 
possesses  great  oxidizing  power. 

HYDROGEN  SULFID  (Hydrosulfuric  Acid),  H2S 
For  the  reactions  use  a  solution  of  H2S  or  (NH4)2S. 

1.  Lead  Acetate    precipitates    black    PbS.      Soluble    in    warm 
dilute  HN03. 

2.  Silver  Nitrate  precipitates  black  Ag2S.     Soluble  in  warm 
dilute  HNO3. 

3.  Antimony  Chlorid  precipitates  orange-red  Sb2S3.     Soluble  in 
(NH4)2S  and  reprecipitated  by  HC1. 

4.  All  soluble  sulfids,  and  most  insoluble  ones,  are  decom- 
posed by  warm  H2SO4,  liberating  H2S,  which  may  be  detected 
by  its  odor ;  also  by  the  brown  or  black  stain  on  a  piece  of  paper 
moistened  with  lead  acetate  and  held  in  the  escaping  gas. 

THIOSULFURIC  ACID,  H2S203 

The  salts  of  this   acid,  which  are  called  thiosulfates,  were 
formerly  called  hyposulfites.     The  free  acid  does  not  exist. 
For  the  reactions  use  a  solution  of  Na2S2O3. 

1.  Lead  Acetate    precipitates    white    PbS2O3.     Soluble    in   an 
excess    of   Na^SgOg.     Decomposed   by  boiling,    forming.  PbS. 
[See  Sulfurous  Acid  1.] 

2.  Silver  Nitrate  precipitates  white  Ag2S2O3.     Easily  soluble 
in  an  excess  of  Na2S2O3,  forming  the  double  salt,  AgNaS2O3. 


REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION     37 

The  precipitate  quickly  beeomes  black,  especially  if  warm,  being 
reduced  to  Ag2S. 

AgaS  A  +  H20  =  Ag2S  +  H2S04. 

3.  Barium  Chlorid  precipitates,  from  a  concentrated  solution, 
white  BaS2O3.     Soluble  in  a  large  quantity  of  water.     Decom- 
posed by  HC1,  liberating  sulfur  dioxid  and  sulfur. 

4.  Ferric  Chlorid  produces  a  violet  color  in  the  solution.     The 
color  is  not  permanent  and  the  solution  soon  becomes  cloudy, 
owing  to  the  reduction  to  FeCl2  and  the  liberation  of  sulfur. 

5.  Hydrochloric   Acid   decomposes    the   thiosulfates,  liberating 
SO2  and  giving  a  precipitate  of  free  sulfur.     [See  Sulfurous 
Acid  4.] 

SULFUROUS  ACID,  H2S03 

This  is  a  weak  acid  and  exists  only  in  a  dilute  solution. 
It  easily  decomposes  when  heated,  forming  SO2  and  H2O.  Its 
salts  are  much  more 'stable,  but  are  all  decomposed  by  dilute 
acids.  A  solution  of  a  sulfite,  on  standing,  becomes  partially 
oxidized,  forming  a  sulfate. 

For  the  reactions  use  a  solution  of  Na2SO3. 

1.  Lead  Acetate  precipitates  white  PbSO3,  which  is  not  decom- 
posed by  boiling.     [See  Thiosulf uric  Acid  1.]     Soluble  in  dilute 
HN03. 

2.  Silver  Nitrate  precipitates  white  Ag2SO3,  which  is  decom- 
posed by  boiling,  forming  black  metallic  silver. 

3.  Barium  Chlorid  precipitates  white  BaSO3.     Easily  soluble 
in  dilute  HC1.     The  solution  in  HC1  is  often  incomplete,  owing 
to  the  presence  of  sulfates,  which  precipitate  insoluble  BaSO4. 

4.  Hydrochloric  Acid  decomposes  the  sulfites,  liberating  SO2. 
If  a  little  potassium  permanganate,  KMnO4,  is  now  added,  it  is 
at  once  decolorized,  owing  to  reduction.     Thus: 

2  KMn04  +  5  S02  +  2  H2O  =  2  MnS04  +  K2S04  +  2  H2S04. 
Most  other  mineral  acids  give  a  similar  reaction. 


38  QUALITATIVE  ANALYSIS 

5.  If  a  solution  of  H2SO3  is  boiled  with  a  little  stannous 
chlorid  and  HC1  it  is  first  reduced  to  H2S,  the  SnCl2  being 
oxidized  to  SnCl4.  The  H2S  then  precipitates  yellow  SnS2. 

SULFURIC  ACID,  H2S04 

For  the  reactions  use  a  solution  of  Na2SO4. 

1.  Lead  Acetate  precipitates  white  PbSO4.     Easily  soluble  in 
ammonium  tartrate  or  ammonium  acetate.      [See  Lead  11.] 

2.  Barium  Chlorid  precipitates  white  BaSO4.     Insoluble  in  all 
dilute  acids. 

3.  Calcium  Chlorid   precipitates,   in   not   too    dilute    solutions, 
white  CaSO4.    A  concentrated  solution  of  Na2CO3  or  (NH4)2CO8 
converts  the  precipitate  into  CaCO3,  which  dissolves  in  dilute 
HC1,  liberating  CO2. 

CHROMIC  ACID,  H2Cr04 

Chromium  ordinarily  acts  like  the  metals,  forming  compounds 
with  the  acid  radicals.  It  may  be  oxidized  by  fusion  with 
Na2CO3  and  KNO3,  forming  a  compound  with  the  metal,  in 
which  chromium  is  found  in  the  acid  radical.  [See  page  15.] 

For  the  reactions  use  a  solution  of  K2CrO4. 

1.  Lead  Acetate  precipitates  yellow  PbCrO4.    Soluble  in  NaOH. 
Insoluble  in  acetic  acid. 

2.  Silver  Nitrate  precipitates  red-brown  Ag2CrO4.     Soluble  in 
HN03  and  in  NH4OH. 

3.  Mercurous  Nitrate  precipitates  brick-red  Hg2CiO4.     Soluble 
in  HNO3. 

4.  Barium  Chlorid  precipitates  yellow  BaCrO4.     Insoluble  in 
NaOH  and  in  acetic  acid. 


REACTIONS  FOR  THE  ACID  RADICALS  IN   SOLUTION     39 

5.  Take  a  dilute  solution  of  hydrogen  dioxid,  H2O2,  acidify 
with  HC1,  add  a  little  ether  (about  half  an  inch  deep  in  the  test- 
tube),  then  two  or  three  drops  of  the  chromate  solution,  and 
shake.    A  portion  of  the  chromate  will  be  oxidized  by  the  H2O2, 
forming  an  unstable  blue  compound,  which  is  supposed  to  be 
perchromic  acid,  HCrO4.     This  dissolves  in  the  ether,  which 
rises  to  the  surface,  giving  it  a  rich  blue  color. 

6.  Nitric  Acid  converts  the  yellow  K2CrO4  into  red  potassium 
dichromate,  K2Cr2O7.     This  salt,  which  may  be  regarded  as  an 
acid  chromate,  gives  in  most  cases  the  same  reactions  as  the 
normal  chromate. 

7.  The    chromates,    and    especially   the    dichromates,    when 
treated  with  H2SO4,  form  sulfates  and  liberate  oxygen.     They 
are  therefore  powerful  oxidizing  agents,  and  are  used  as  such, 

especially  in  organic  chemistry. 

k      (\   v  o       -1-  '•))!  SOi,    -*    K^Ou  ^^i  r^V).  *vAfcG*-»3  ^J 

'  2  ^       -j  '  '  '>         ^  *•  X> 

NITROUS  ACID,  HN02 

This  acid  does  not  exist  in  the  free  state.  Even  when  liber- 
ated in  a  dilute  solution  it  is  easily  decomposed,  giving  nitric 
acid,  nitric  oxid,  and  water.  Its  salts,  the  nitrites,  are  quite 
stable,  but  they  are  all  decomposed  by  dilute  acids,  forming 
HNO3  and  liberating  NO. 

For  the  reactions  use  a  solution  of  KNO2. 

1.  Silver  Nitrate  precipitates,  in  a  concentrated  solution,  white 
AgN02. 

2.  Cobalt  Nitrate,  to  which  has  been  added  acetic  acid,  precipi- 
tates in  an  excess  of   the  solution,  yellow  Co(NO2)8(KNO2)3. 
[See  Cobalt  10.] 

Sodium  nitrite  does  not  give  this  reaction. 

3.  If  KI  is  added  to  a  solution  of  a  nitrite,  together  with  a 
little  starch  paste  and  a  few  drops  of  dilute  H2SO4,  iodin  is 
liberated,  which  colors  the  starch  paste  blue. 


40  QUALITATIVE   ANALYSIS 

4.  If  a  little  FeSO4  is  added  to  a  solution  of  a  nitrite,  and 
then  a  few  drops  of  dilute  acetic  acid,  the  whole  becomes  brown, 
from  the  NO  which  is  liberated  dissolving  in  the  FeSO4.     [See 
Nitric  Acid  1.] 

5.  Nitrous  acid  is  capable  of  oxidation  to  nitric  acid  in  the 
presence  of  oxidizing  agents.     If  KMnO4  is  added  to  a  solution 
of  a  nitrite  acidified  with  H2SO4,  it  is  decolorized,  owing  to  its 
reduction.     Thus : 

4  KMn04  +  10  KNO2  +  11  H2S04  = 

7  K2S04  +  4  MnSO,  +  10  HN03  +  6  H20. 

NITRIC  ACID,  HN03 

All  nitrates  are  soluble  in  water,  and  so  form  no  precipitates 
with  the  metals.  They  are  all  decomposed  by  H2SO4,  liberating 
HNO3.  Nitric  acid  is  an  oxidizing  agent,  and  the  tests  which 
indicate  its  presence  are  connected  with  an  oxidizing  action. 

For  the  reactions  use  a  solution  of  KNO3. 

1.  Mix  some  of  the  nitrate  solution  with  an  equal  volume  of 
FeSO4.     Incline  the  tube  a  little  and  carefully  pour  down  the 
side  some  concentrated  H2SO4,  and  where  the  mixture  meets 
the  surface  of  the  acid  a  brown  ring  of  color  will  appear.     The 
brown  compound,  which  is  due  to  a  solution  of  NO  in  FeSO4, 
is  decomposed  by  heat,  liberating  the  NO.     [See  Nitrous  Acid  4.] 

The  action  is  threefold:  (1)  the  liberation  of  HNO3  by  the 
action  of  H2SO4  on  the  nitrate ;  (2)  the  oxidation  of  FeSO4  by 
the  HNO3,  liberating  NO ;  and  (3)  the  absorption  of  NO  by 
FeSO4,  forming  the  unstable  brown  compound. 

2.  If  concentrated  H2SO4  is  mixed  with  the  nitrate  solution, 
a  few  fragments  of  copper  added,  and  the  whole  boiled,  NO  is 
liberated,  which,  combining  with  the  oxygen  in  the  air,  forms 
red-brown  fumes  of  NO2.     These  will  be  more  easily  seen  by 
looking  down  through  the  mouth  of  the  tube.     The  action  is 
analogous  to  that  in  the  first  test. 


REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION     41 

PHOSPHORIC  ACID,  H3P04 
For  the  reactions  use  a  solution  of  Na2HPO4. 

1.  Lead  Acetate  precipitates  white  Pb3(PO4)2.     Easily  soluble 
in  HNO3. 

2.  Silver    Nitrate    precipitates    yellow    Ag3PO4.     Soluble    in 
N1I4OH  and  in  HNO3. 

3.  Barium  Chlorid   precipitates   white    BaHPO4.     Soluble    in 
dilute  HC1  and  reprecipitated  by  NH4OH. 

4.  Magnesium  Sulfate,  to  which  has  been  added  NH4C1  and  then 
NH4OH  in  excess,  precipitates  white  crystalline  MgNH4PO4. 
This  precipitate  forms  slowly,  and  in  a  dilute  solution  is  com- 
plete only  after  standing  some  hours. 

5.  Ammonium    Molybdate,    (NH4)2MoO4,    with    an   'excess    of 
HNO3    precipitates     yellow     ammonium     phospho-molybdate, 
(NH4)3PO4(MoO3)12.     Soluble    in    NH4OH    and    in    excess    of 
Na2HPO4.     The  precipitation  is  hastened  by  warming. 

2     4.0   = 
ARSENIOUS  ACID,  H3As03 

We  have  already  learned  that  arsenic  may  be  precipitated  as 
a  sulfid.  [See  Arsenic  1.]  It  forms  no  other  salts  in  which 
it  acts  as  a  metal,  but  acts  like  an  acid,  forming  salts  with  the 
metals.  The  alkaline  salts  only  are  soluble  in  water.  These 
may  be  formed  by  dissolving  the  oxid,  As2O3,  in  a  solution  of 
an  alkaline  hydroxid. 

For  the  reactions  use  a  solution  of  K3AsO3  or  Na3AsO3. 

1.  Silver  Nitrate  precipitates,  in  a  neutral  solution,  yellow 
Ag3AsO3.  Soluble  in  NH4OH,  in  HNO3,  and  in  NH4NO3.  If 
the  ammoniacal  solution  is  boiled  for  some  time,  metallic  silver 
is  precipitated,  a  portion  of  the  arsenite  being  oxidized  to  an 
arsenate. 


42  QUALITATIVE   ANALYSIS 

2.  Copper    Sulfate    precipitates    Scheele's    green,    CuHAsO3. 
Soluble  in  NH4OH  and  in  acids. 

If  an  excess  of  NaOH  is  added  to  a  solution  of  an  arsenite, 
then  a  few  drops  of  CuSO4,  and  the  whole  boiled,  red  Cu2O 
is  precipitated,  a  portion  of  the  arsenite  being  oxidized  to  an 
arsenate.  (Distinction  from  arsenates.)  [See  Arsenic  Acid  2.] 

3.  Reinsch's  Test.     If  a  piece  of  metallic  copper  is  placed  in 
an  arsenite  solution  acidified  with  HC1  and  warmed,  a  gray 
film  of  copper  arsenid,  Cu5As2,  is  formed  on  the  surface  of  the 
copper. 

4.  Stannous  Chlorid,   to   which  has  been   added  at  least  two 
volumes  of  concentrated  HC1,  and  then  warmed,  precipitates 
black  metallic  arsenic. 

5.  Hydrogen  or  Ammonium  Sulfid  gives  no  precipitate  in  a  neutral 
or  alkaline  solution  of  an  arsenite.     If  the  solution  is  acid,  or  is 
rendered  acid,  yellow  As2S3  is  precipitated.     [See  Arsenic  1.] 

ARSENIC  ACID,  H3As04 

Salts  of  this  acid  are  derived  from  As2O5  or  by  oxidation  of 
the  arsenites. 

The  reactions  for  arsenic  acid  are  very  closely  analogous  to 
those  for  phosphoric  acid. 

For  the  reactions  use  a  solution  of  Na3AsO4. 

1.  Silver  Nitrate  precipitates,  in  a  neutral  solution,  red-brown 
Ag3AsO4.     Soluble  in  NH4OH  and  in  HNO3. 

2.  Copper  Sulfate  precipitates  light  blue  CuHAsO4.     Soluble 
in  NH4OH.     Not  decomposed  by  boiling  in  an  alkaline  solution. 
[See  Arsenious  Acid  2.] 

3.  Magnesium  Sulfate,  to  which  has  been  added  NH4C1  and  then 
NH4OH  in  excess,  precipitates  white  crystalline  MgNH4AsO4. 
The  precipitate  forms  slowly,  especially  in  a  dilute  solution,  but 
may  be  hastened  by  shaking.     Soluble  in  HC1  and  reprecipitated 


REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION     43 

by  NH4OH.  If  (NH4)2S  or  H2S  is  added  to  the  solution  in 
HC1,  the  arsenate  is  reduced  and  yellow  As2S3  mixed  with  free 
sulfur  is  precipitated.  The  action  is  hastened  by  warming. 
(Distinction  from  phosphates.) 

4.  Ammonium  Molybdate  gives  no  precipitate  in  the  cold.     If 
the   solution  is  warmed,  yellow  ammonium  arseno-molybdate, 
(NH4)3AsO4(MoO3)12,  is  formed.     Insoluble  in  NH4OH. 

5.  Hydrogen  or  Ammonium  Sulfid  in  a  solution   acidified  with 
HC1  reduces  an  arsenate  to  an  arsenite,  and  then  precipitates 
yellow  As2S3  mixed  with  free  sulfur.     [See  Arsenic  1.] 

BORIC  ACID,  H3B03 

This  is  a  weak  acid.  Only  a  few  salts  are  known,  and  most 
of  these  are  from  the  derivatives, — metaboric  acid,  HBO2,  and 
pyroboric  acid,  H2B4O7. 

For  the  reactions  use  a  solution  of  borax,  Na2B4O7. 

1.  Lead  Acetate  precipitates,  in  not  too  dilute  solutions,  white. 
Pb(BO2)2.     Soluble  in  excess  of  the  reagent. 

2.  Silver  Nitrate  precipitates  white  AgBO2.     If  the  solution  is 
too  dilute,  the  precipitate  is  yellow  or  brown,  from  the  presence 
of  Ag20. 

3.  Barium  Chlorid  precipitates,  in  not  too  dilute  solutions,  white 
Ba(BO2)2.     Soluble  in  excess  of  the  reagent  and  in  NH4C1. 

4.  If  a  little  alcohol  is  added  to  a  solution  of  a  borate  in  a 
small  porcelain  dish,  then  a  little  concentrated  H2SO4,  and  the 
whole  warmed,  ethyl  borate  (boric  ether)  is  formed.     This   is 
inflammable,  and  so  the  mixture  may  be  ignited,  when  it  will 
be  seen  to  burn  with  a  green-bordered  flame. 

CARBONIC  ACID,  H2C03 

This  is  a  weak  acid,  existing  only  in  a  dilute  solution.  Its 
salts  are  common  and  quite  stable. 

For  the  reactions  use  a  solution  of  Na,2CO3. 


44  QUALITATIVE   ANALYSIS 

1.  Lead  Acetate  precipitates  white  PbCO3  or,  if  the  solution 
is  hot,  a  white  basic  carbonate. 

2.  Silver  Nitrate  precipitates  light  yellow  Ag2CO3.     Soluble 
in  NH4OH  and  in  (NH4)2CO3. 

3.  Barium  Chlorid  precipitates  white  BaCO3.     Soluble  in  water 
containing   CO2,   forming   BaH2(CO3)2,  and  reprecipitated   by 
boiling.     To  show  this  solubility  fill  a  large  test-tube  containing 
the  precipitate  with  water,  shake  the  mixture,  pour  out  about 
one  half,  and  lead  CO2  through  the  liquid  until  the  precipitate 
dissolves. 

4.  All  carbonates,  whether  soluble  or  insoluble  in  water,  are 
decomposed  with  effervescence  by  dilute  HC1  and  other  acids, 
liberating  CO2.     The  presence  of  the  CO2  may  be  confirmed 
by  decanting  the  heavy  gas  into  a  test-tube  containing  a  little 
lime-water,  Ca(OH)2.     On  shaking  with  the  gas,  the  lime-water 
becomes  milky,  owing  to  the  precipitation  of  CaCO3.     Baryta- 
water,  Ba(OH)2,  may  be  used  in  place  of  the  lime-water. 


SILICIC  ACID,  H4Si04 

This  acid  exists  only  in  a  dilute  solution.  On  attempting  to 
concentrate  the  solution,  it  loses  a  molecule  of  water  and  forms 
metasilicic  acid,  H2SiO3.  This  in  turn  decomposes  on  heating 
to  130°,  forming  H2O  and  SiO2.  The  silicates  occurring  in 
nature  are  generally  salts  of  the  polysilicic  acids.  Only  the 
alkaline  silicates  are  soluble  in  water. 

For  the  reactions  use  a  solution  of  Na4SiO4. 

1.  Hydrochloric  Acid  (concentrated)  precipitates,  in  not  too  dilute 
solutions,  white  gelatinous  H4SiO4.  In  dilute  solutions  this 
precipitate  will  appear  only  after  long  standing.  It  is  somewhat 
soluble  in  water  and  in  HC1,  so  that  a  portion  of  it  remains  in 
the  solution. 


REACTIONS  FOR  THE  ACID  RADICALS   IN  SOLUTION     45 
2.    Ammonium   Carbonate  or  Ammonium  Chlorid   precipitates   the 


same. 


3.  Barium  Chlorid  precipitates  white  Ba2SiO4. 

4.  Calcium  Chlorid  precipitates  white  Ca2SiO4 


ACETIC   ACID,  H(C2H302) 

This  and  the  following  acids  belong  to  the  division  known 
as  Organic  Chemistry.  They  are,  however,  frequently  used  in 
the  process  of  analysis,  and  it  is  therefore  desirable  to  know 
how  to  recognize  them. 

Acetic  acid  is  a  liquid,  boiling,  when  pure,  at  119°.  It  has 
a  characteristic  pungent  odor  like  that  of  vinegar.  It  is  a 
monobasic  acid,  and  its  salts  are  all  soluble  in  water,  although 
one  or  two  of  them  may  be  partly  precipitated  if  the  solutions 
are  sufficiently  concentrated. 

For  the  reactions  use  a  solution  of  Na(C2H3O2). 

1.  Silver  Nitrate  precipitates,  from  concentrated  solutions,  white 
Ag(C2H302).     Soluble  in  NH4OH. 

2.  Ferric  Chlorid  gives,  in  a  neutral  solution,  a  deep  red  colora- 
tion, but  no  precipitate.      If  this  solution  is  now  boiled,  a  pre- 
cipitate  of   red-brown   basic   ferric    acetate,    Fe(OH)2(C2H3O2), 
appears  and  the  solution  becomes  colorless. 

3.  Sulfuric  Acid  gives  no  precipitate,  but  liberates  acetic  acid, 
which  may  be  recognized  by  its  odor. 

4.  If  a  few  drops  of  alcohol  are  added  to  a  solution  of  an 
acetate,  then  a  little  concentrated  sulfuric  acid,  and  the  whole 
warmed,  ethyl  acetate  (acetic  ether)  is  formed.     This  is  a  vola- 
tile, ethereal  liquid,  having  an  agreeable  and  characteristic  odor 
somewhat  like  that  of  apples. 


46  QUALITATIVE  ANALYSIS 

OXALIC  ACID,  H2C204 

Oxalic  acid  is  a  white  crystalline  solid,  soluble  in  water.  It 
is  a  dibasic  acid.  The  oxalates  of  the  alkalies  are  soluble  in 
water,  while  most  of  the  others  are  insoluble. 

For  the  reactions  use  a  solution  of  (NH4)2C2O4. 

1.  Lead  Acetate  precipitates  white  PbC2O4.     Soluble  in  HNO3. 

2.  Silver    Nitrate    precipitates    white    Ag2C2O4.     Soluble    in 
HN03  and  in  NH4OH. 

3.  Barium  Chlorid  precipitates,  in  not  too  dilute  solutions,  white 
BaC2O4.     Somewhat  soluble  in  water.     Soluble  in  HC1  and  in 
acetic  acid. 

4.  Calcium  Chlorid   precipitates  white   CaC2O4.     Insoluble  in 
water,  in  acetic  acid,  and  in  NH4OH.     Soluble  in  HC1. 

TARTARIC  ACID,  H2(C4H406) 

Tartaric  acid  is  a  white  crystalline  solid,  soluble  in  water. 
It  is  a  dibasic  acid.  The  normal  tartrates  of  the  alkalies  are 
easily  soluble  in  water,  while  the  acid  salts  dissolve  with 
difficulty.  Most  of  the  other  tartrates  are  insoluble  in  water, 
but  many  of  them  dissolve  in  excess  of  the  alkaline  tartrates, 
forming  double  salts. 

For  the  reactions  use  a  solution  of  sodium  potassium  tartrate 
(Rochelle  salt),  %aK(C4H4O6). 

1.  Lead  Acetate  precipitates  white  Pb(C4H4O6).     Soluble  in 
HNO3  and  in  NH4OH. 

2.  Silver  Nitrate  precipitates  white^  Ag2(C4H4O6).     Soluble  in 
HNp3am\inNH4OH. 

3.  Barium  Chlorid,  when  added  in  excess,  precipitates  white 
Ba(C4H4O6).     Soluble  in  acetic  acid. 


REACTIONS  FOR  THE  ACID  RADICALS  IN  SOLUTION     47 

4.  Calcium  Chlorid,  when  added  in  excess,  precipitates  white 
Ca(C4H4O6).     Soluble  in  all  acids  and  in  NH4C1.     Insoluble  in 
NH4OH. 

5.  If  a  few  drops  of  silver  nitrate  are  added  to  a  solution  of 
a  tartrate  in  a  carefully  cleaned  test-tube,  then  NH4OH  added 
drop  by  drop  until  the  precipitate  first  formed  is  nearly  dis- 
solved, and  the  whole  gently  warmed,  the  silver  tartrate  will  be 
reduced  and  metallic  silver  deposited  as  a  brilliant  mirror  upon 
the  glass.     (A  very  characteristic  reaction.) 


PART  II 


REACTIONS    FOR   DRY    SUBSTANCES 

BLOWPIPE  ANALYSIS 

Matter,  at  the  ordinary  temperature,  nearly  always  exists  in 
the  solid  state.  For  purposes  of  analysis  it  is  much  more  con- 
venient to  have  it  in  the  liquid  condition.  This  may  be  obtained 
by  solution  or  fusion.  Since  the  latter  often  requires  a  very 
high  temperature  the  former  is  almost  universally  employed. 

In  obtaining  a  solution  of  a  given  substance,  it  sometimes 
happens  that  a  change  in  composition  occurs,  since  solution 
may  be  attended  by  chemical  change.  In  order  to  know  the 
original  composition  of  a  substance  we  may  have  to  analyze  it 
in  its  original  condition.  The  methods  employed  for  this  pur- 
pose belong  to  that  part  of  analytical  chemistry  known  as 
Blowpipe  Analysis,  and,  while  it  is  not  desirable  to  present  an 
exhaustive  treatise  at  this  time,  a  knowledge  of  the  simpler 
operations  belonging  to  this  part  of  the  subject  is  indispensable 
to  the  chemist. 

By  these  methods  all  simple  inorganic  substances  and  many 
organic  compounds  may  be  completely  analyzed,  and  many 
important  facts  about  the  more  complex  compounds  may  be 
learned.  Some  of  the  phenomena,  while  perfectly  evident,  are 
so  slight  as  to  be  easily  overlooked,  unless  carefully  observed. 
They  are,  however,  simple  and  easy  to  follow,  and  if  carefully 
observed  the  results  are  accurate  and  conclusive.  This  part  of 
the  work,  therefore,  is  of  great  value  in  developing  the  powers 
of  observation. 

48 


REACTIONS  FOR  DRY   SUBSTANCES  49 

Nearly  all  of  the  results  are  produced  by  heat,  either  alone  or 
with  reagents.  The  operations  should  be  carried  on  system- 
atically and  the  student  taught  to  make  the  proper  deductions 
from  each  operation  before  going  on  to  the  next. 

For  the  heat  a  Bunsen  lamp  is  employed.  This  is  supple- 
mented by  an  instrument  called  a  blowpipe.  The  lamp  should 
be  furnished  with  an  inner  tube  to  be  inserted  for  use  with  the 
blowpipe.  The  latter  is  used  for  producing  both  the  oxidizing 
and  reducing  flame.  The  proper  manner  of  using  the  blow- 
pipe, and  the  ability  to  produce  both  the  oxidizing  and  reducing 
flame,  should  be  thoroughly  mastered. 

The  operation,  which  is  not  very  difficult,  but  is  apt  to  puzzle 
the  student  at  first,  is  as  follows.  Insert  the  blowpipe  tube 
in  the  Bunsen  lamp  and  turn  down  the  flame  so  that  it  will  be 
from  four  to  live  centimeters  in  length.  Bring  the  tip  of  the 
blowpipe  into  the  flame  about  a  third  part  of  the  width  of  the 
flame  and  near  the  end  of  the  inner  tube.  If  a  fairly  strong 
current  of  air  is  now  sent  through  the  blowpipe,  a  long,  pointed, 
blue,  oxidizing  flame  is  produced.  This  flame  is  used  for  fusion 
and  oxidation.  The  hottest  part  of  the  flame  is  about  midway 
between  the  point  of  the  inner  blue  cone  and  the  extreme  tip 
of  the  flame.  The  point  of  maximum  oxidation  is  at  the  extreme 
tip  of  the  flame,  or  even  just  beyond  this  if  the  temperature  is 
found  to  be  sufficiently  high. 

If  the  tip  of  the  blowpipe  is  held  just  outside  the  gas  flame, 
and  a  gentle  stream  of  air  is  sent  through  the  blowpipe,  the 
inner  blue  cone  will  be  surrounded  by  a  luminous  mantle,  form- 
ing the  reducing  flame.  The  point  of  maximum  reduction  is 
just  within  the  point  of  the  luminous  mantle.  The  reducing 
flame  is  not  nearly  so  hot  as  the  oxidizing  flame. 

The  proper  use  of  the  blowpipe  can  best  be  learned  by  prac- 
tice under  the  guidance  of  a  competent  instructor.  The  other 
apparatus,  as  well  as  the  reagents  to  be  used,  will  be  described 
as  they  are  employed. 


50  QUALITATIVE   ANALYSIS 

I.  THE  EFFECT  OF  HEAT  ALONE 

The  substance  is  heated  in  a  piece  of  hard  glass  tubing  closed 
at  one  end.  It  should  be  heated  in  the  Bunsen  flame,  at  first 
gently,  then  strongly. 

A.  Water  is  given  off.  This  is  recognized  by  its  condensation 
in  small  drops  in  the  upper  part  of  the  tube,  and  indicates  the 
following  about  the  substance. 

(a)  It  is  a  deliquescent  salt.     A  deliquescent  substance  is  one 
which  absorbs  moisture  from  the  atmosphere.     The  water  is 
driven  off  at  a  comparatively  low  temperature,   and   usually 
in  small  quantities. 

(b)  It  contains  enclosed  water.     The  substance  often  decrepi- 
tates, or  crackles,  when  heated.     This  is  caused  by  the  bursting 
of  the  particles.     The  amount  of  water  is  usually  small. 

This  reaction  is  particularly  characteristic  of  NaCl  and 
certain  other  halogen  salts ;  also  of  other  salts  which  do  not 
contain  water  of  crystallization. 

(c)  It    contains   chemically  combined  water.     The  substance 
may  be  an  hydroxid,  an  acid  salt  (generally  of  a  volatile  acid  or 
one  easily  decomposed),  an  ammonium  salt,  or  an  organic  com- 
pound.    The  steam,  as  it  passes  out  of  the  tube,  should  be 
tested  with  a  piece  of  moistened  litmus  paper.     A  neutral  reac- 
tion usually  indicates  an  hydroxid;   an  acid  reaction,  an  acid 
salt;  an  alkaline  reaction,  an  ammonium  salt.     The  compound 
often  shows  a  permanent  change  of  color.     If  it  blackens  and 
gives  off  empyreumatic  odors,  it  indicates  an  organic  substance. 

This  action  may  require  a  fairly  high  temperature.  The 
amount  of  water  varies,  but  is  usually  not  very  great. 

(d)  It  contains  water   of  crystallization.     A  portion   of  the 
water  comes  off  at  or  below  100°,  but  the  last  portion  may 
require  a  much  higher  temperature.     The  amount  of  water  is 


REACTIONS  FOR  DRY  SUBSTANCES  51 

usually  relatively  large.     Some  alums,  borates,  and  phosphates 
swell  up  considerably  while  giving  off  their  water. 

B.  A  gas  is  given  off.  This  may  consist  of  one  or  more  of  the 
following  gases. 

(a)  Oxygen.     This  is  recognized  by  the  igniting  of  a  glow- 
ing   splinter  when    introduced    into    the    tube.     It    indicates 
that  the  substance  was  a  nitrate,  a  peroxid,  or  some  highly 
oxidized  salt,  such  as  a  chlorate,  bromate,  iodate,  dichromate, 
or  permanganate. 

(b)  Ammonia.     This  is  easily  recognized  by  its  odor  and  the 
white   fumes   of  NH4C1  which  are  formed  when  a  glass  rod 
moistened  with  HC1  is  held  in  the  escaping  gas.     It  indicates 
an  ammonium  salt. 

(c)  Carbon  dioxid.    This  is  recognized  by  the  turbidity  which 
is  caused  when  lime-water  [Ca(OH)2]  is  exposed  to  the  gas. 
It  indicates  a  carbonate  or  an  organic  compound.     The  latter 
usually  blackens  by  heating.     All  carbonates  give  this  except 
normal  carbonates  of  the  alkali  metals.     These,  especially  the 
commercial  carbonates,  sometimes  contain,  as  impurity,  small 
quantities  of  the  acid  carbonates,  which  give  this  reaction.     A 
convenient  instrument  for  showing  this  is  made  from  a  piece 
of  glass  tubing  drawn  out  at  one  end  to  a  capillary  and  bent 
like  a  siphon.     This  is  attached  to  the  closed  tube  by  a  piece  of 
rubber  tubing,  the  capillary  end  being  placed  in  the  lime-water 
and  the  closed  tube  heated. 

(d)  Carbon  monoxid.     This    gas    burns  with  a  bright  blue 
flame,  by  which  it  may  be  recognized.    This  indicates  an  oxalate 
or  a  formate.     The  latter  blackens  when  heated. 

The  flame  of  carbon  monoxid  does  not  always  appear  blue, 
because  of  the  presence  of  some  impurity  (such  as  sodium). 
It  may  also  fail  to  appear  because  of  the  presence  of  water  in 
the  form  of  steam. 


52  QUALITATIVE   ANALYSIS 

(e)  Sulfur  dioxid.     This  may  be  recognized  by  its  odor,  that 
of  burning  sulfur.     It  indicates  a  sulfate  or  a  sulfite. 

(f)  Hydrogen  sulfid.     This   is   recognized  by  its  odor,  and 
by  blackening  a  piece  of  paper  moistened  with  lead  acetate.     It 
indicates  a  sulfid  containing  water. 

(g)  Nitrous  oxid.     This  supports  combustion  nearly  as  well  as 
oxygen  and  may  be  recognized  by  the  same  test.     It  indicates 
ammonium  nitrate. 

(h)  Nitrogen  trioxid  or  tetroxid.  These  gases  may  be  recog- 
nized by  their  brownish-red  color  and  a  peculiar  odor  which 
is  like  that  of  nitric  acid.  They  usually  indicate  a  nitrate  or 
nitrite  of  the  heavy  metals.  The  alkaline  salts  do  not  give 
this  reaction. 

(i)  Chlorin.  This  is  recognized  by  its  yellow  color  and  its 
odor;  also  by  its  bleaching  action.  It  indicates  certain  chlorids 
and  hypochlorites. 

(/)  Bromin.  This  is  recognized  by  its  red-brown  color  and 
its  odor,  which  is  much  like  that  of  chlorin.  It  indicates  certain 
bromids,  and  other  bromin  compounds. 

(k)  lodin.  This  is  recognized  by  its  deep  violet-colored  vapor. 
It  indicates  iodin,  an  iodid,  or  some  other  iodin  compound. 

(I)  Cyanogen.  This  is  recognized  by  its  odor  (Poison !), 
which  is  like  that  of  KCN,  and  by  the  crimson  color  of  its  flame. 
It  indicates  a  cyanid  of  one  of  the  less  basic  metals. 

(m)  Organic  gases.  These  may  usually  be  recognized  by 
their  inflammability,  the  flame  being  more  or  less  luminous. 
They  indicate  an  organic  substance. 

C.  A  sublimate  is  formed.  Some  substances  when  heated  pass 
directly  from  the  solid  to  the  gaseous  state.  When  this  gas 
comes  in  contact  with  the  colder  surface  of  the  upper  part  of  the 
tube  it  condenses  again  to  a  solid,  forming  a  sublimate.  This 


REACTIONS   FOR  DRY  SUBSTANCES  53 

usually  begins  to  form  at  a  distance  of  one  or  more  centimeters 
from  the  substance,  and  the  line  of  formation  is  usually  a  sharp 
one,  the  sublimate  shading  off  gradually  above  it.  Sometimes 
\\hen  a  substance  fuses,  a  film  of  the  melted  material  may 
extend  up  the  tube  for  some  distance  from  the  substance.  This 
must  not  be  mistaken  for  the  sublimate.  The  sublimates  vary 
much  in  color. 

1.  A  white  sublimate  is  formed  by  the  following  substances. 

(a)  Ammonium  salts.     If  two  or  three  drops  of  a  solution  of 
NaOH  are  placed  in  the  tube  with  the  substance  and  warmed 
(the  original  substance  may  be  treated  this  way  in  a  test-tube), 
ammonia  is  given  off,  which  may  be  recognized  by  its  odor. 

(b)  Mercurous  chlorid.     The  sublimate  is  yellow  while  hot, 
but  becomes  white  on  cooling. 

(c)  Mercuric    chlorid.     This    is    much    like    the    mercurous 
chlorid,  but  fuses  before  it  sublimes. 

(d)  Arsenic  trioxid.     This  gives   a  sublimate   of  octahedral 
crystals.     If  a  bit  of  charcoal  is  placed  in  the  tube  with  the 
substance  and  heated,  a  black  mirror  of  arsenic  is  produced. 

2.  A  colored  sublimate  is  formed  by  the  following  substances. 

(a)  Arsenic.     This  gives  a  black  shining  mirror,  and  may  be 
formed  by  the  element  itself,  or  by  its  compounds  in  the  pres- 
ence of  a  reducing  agent,  like  carbon.    The  vapor  has  a  peculiar 
garlic-like  odor. 

(b)  Antimony    sulfid.     This    sublimes    only  at  a  very  high 
temperature,  the  sublimate  being  black  when  hot,  and  reddish 
brown  when  cold. 

(c)  Mercuric   sulfid.     This    forms  a   black  sublimate  which 
shows  red  when  rubbed  with  a  glass  rod. 

(d)  lodin.     This  gives  a  black  sublimate  and  a  deep  violet- 
colored  vapor. 


54  QUALITATIVE   ANALYSIS 

(e)  Arsenic  sulfid.  This  is  dark  reddish  brown  while  hot, 
and  yellowish  red  when  cold. 

(/)  Sulfur.  This  may  come  from  free  sulphur  or  from  cer- 
tain sulfids.  The  sublimate  is  brownish  yellow  while  hot,  and 
sulfur  yellow  when  cool.  It  burns  easily,  giving  off  sulfur 
dioxid. 

(g)  Mercuric  iodid.  This  forms  a  yellow  sublimate,  which 
soon  changes  to  red,  especially  if  rubbed  with  a  glass  rod. 

D.  The  substance  changes  color.  Many  substances,  upon  heating, 
do  not  change  in  composition,  but  change  in  appearance.  Upon 
cooling,  the  original  color  appears. 

Many  salts  decompose  on  heating,  leaving  an  oxid  of  the 
metal  which  may  have  a  bright  color.  This  may  exhibit  some 
of  the  phenomena  given  below,  and  thus  give  a  clew  to  the 
original  substance.  The  following  are  characteristic. 

(a)  The  substance  is  white,  becomes  yellow  when  hot,  and 
white  again  on  cooling.     This  indicates  zinc  oxid,  ZnO. 

(b)  The  substance  is  white,  becomes  yellowish  brown  when  hot, 
and  is  a  dirty  pale  yellow  on  cooling.     At  a  high  temperature 
it  is  infusible  and  luminous.     It  indicates  stannic  oxid,  SnO2. 

(c)  The  substance  is  orange  or  light  yellow,  becomes  brown 
red  when  hot,  and  yellow  when  cold.     It  is  fusible  at  a  high 
temperature.     It  indicates  lead  oxid,  PbO. 

(d)  The  substance  is  yellow,  becomes  orange  yellow  or  red 
brown  when  hot,  and  yellow  when  cold.     It  is  fusible  at  a  high 
temperature.     It  indicates  bismuth  oxid,  Bi2O3. 

(e)  The  substance  is  red  or  red  brown,  becomes  a  very  dark 
red  brown,  almost  black,  when  hot,  and  red  when  cold.     It  is 
infusible.     It  indicates  ferric  oxid,  Fe2O3. 

(/)  The  substance  is  yellowish  red,  becomes  dark  brown  or 
black  when  hot,  and  red  when  cold.  It  decomposes  when 


REACTIONS   FOR  DRY  SUBSTANCES  55 

strongly  heated,  forming  a  black  or  gray  sublimate,  and  giving 
off  oxygen.     It  indicates  mercuric  oxid,  HgO. 

E.  The  substance  fuses  without  decomposition.     This  generally 
indicates  that  the  substance  is  an  alkaline  salt,  though  a  very 
few  other  salts  do  this. 

F.  The   substance  carbonizes.     Water  is  usually  given  off  to- 
gether with  gases  having  a  characteristic  odor.     This  indicates 
an  organic  compound.     Only  a  few  organic  compounds  can  be 
completely  determined  by  these  actions.     They  are  the  more 
common  acids,  and  metallic  salts  of  organic  acids. 

(a)  Acetates.     These  give  off  aceton,  which  has  a  characteris- 
tic odor  somewhat  suggestive  of  vinegar. 

(b)  Formates.     These  give  off  CO,  which  burns  with  a  bright 
blue  flame. 

(c)  Tartrates.     These  give  an  odor  like  that  of  burnt  sugar. 

These  salts  of  metallic  bases  and  organic  acids  in  decompos- 
ing by  heat  always  leave  a  carbonate  of  the  metal,  which  by 
further  heating  will  form  an  oxid  of  the  metal,  unless  the  latter 
belongs  to  the  alkali  group.  The  carbonate  may  be  recognized 
by  the  effervescence  when  a  drop  of  HC1  is  added. 

H.  THE  SUBSTANCE  IS  HEATED  ON  CHARCOAL 

This  gives  the  effect  of  heat  in  the  presence  of  a  strong 
reducing  agent,  the  hot  charcoal.  A  shallow  cavity  is  made 
in  a  piece  of  soft-wood  charcoal,  the  substance  is  placed  in  this, 
and  heated  with  the  blowpipe  flame,  at  first  gently,  afterwards 
strongly.  It  is  best  to  hold  the  charcoal  somewhat  inclined 
toward  the  flame,  so  that  in  case  an  incrustation  should  be 
formed  it  may .  be  observed  more  readily.  If  the  substance 
is  very  light  and  dry,  so  that  it  is  liable  to  be  blown  away,  it 
may  be  moistened  with  water,  or,  in  some  cases,  a  small  piece 
of  borax  may  be  fused  with  the  substance. 


56  QUALITATIVE   ANALYSIS 

A.  The  substance  fuses  easily  without  decomposition,  and  sinks 
into  the  charcoal.     This  indicates  a  salt  of  the  alkali  metals  or 
some  of  the  salts  of  the  alkaline  earths. 

B.  The  substance  yields  a  metallic  bead  without  any  incrustation. 

(a)  The  bead  is  white.    This  indicates  tin,  aluminum,  or  silver. 
Tin  is  very  easily  fusible  (232°),  aluminum  requires  quite  a 
high    temperature    (655°),    and    silver    fuses    only   with   great 
difficulty  (960°). 

(b)  The  bead  is  red.     This  indicates  copper  and  requires  a 
very  high  temperature  (1080°)  and  often  long-continued  heating. 
The  bead  is  malleable,  which  distinguishes  it  from  CuO,  which 
is  red,  but  brittle. 

(c)  The  bead  is  yellow.     This  indicates  gold  and  requires  a 
very  high  temperature  (1061°). 

C.  The  substance  yields  a  metallic  bead  with  an  incrustation. 

(a)  The  bead  is  white,  soft,  and  malleable.     The  incrustation 
is  yellow  and  volatile.     This  indicates  lead. 

(b)  The  bead  is  white,  soft,  and  malleable.     The  incrustation 
is  red  brown  and  volatile.     This  indicates  cadmium. 

(c)  The   bead  is  white,  rather  hard,  and  malleable.      It  is 
pretty  well    covered  with  the  incrustation,   which   is   yellow 
when  hot,  and  white  when  cold.      This  indicates  zinc. 

(d)  The  bead  is  white,  hard,  and  brittle.     The  incrustation 
is  white  and  volatile.     This  indicates  antimony. 

(e)  The  bead  is  white,  hard,  and  brittle.     The  incrustation 
is  yellow  and  volatile.     This  indicates  bismuth. 

D.  The  substance  is  infusible,  dark  brown  or  black  in  color,  gives 
no  incrustation,  and  is  more  or  less  easily  attracted  by  a  magnet. 

This  indicates  iron,  chromium,  nickel,  cobalt,  or  manganese. 
These  may  be  distinguished  by  the  borax  bead.  [See  Coloration 
of  the  Borax  or  Microcosmic  Bead.] 


REACTIONS   FOR   DRY   SUBSTANCES  57 

Molybdenum,  tungsten,  and  some  of  the  platinum  metals  are 
infusible,  and  would  be  found  here.     They  are  not  magnetic. 

E.  The  substance  deflagrates,  or  burns  up  quickly.      This  indi- 
cates a  nitrate,  or  some  highly  oxidized  salt,  such  as  a  chlorate, 
bromate,  or  iodate.     Only  the  alkaline  salts  show  this  action  in 
a  marked  degree ;  other  salts  are  scarcely  to  be  recognized  by 
this  test. 

F.  The  substance  decrepitates.      This  indicates  a  crystalline  salt, 
which  may  contain  enclosed  water  but  does  not  usually  contain 
water   of    crystallization.      Sodium  chlorid  and   other  halogen 
salts  show  this  action  best. 

G.  The  substance  volatilizes. 

(a)  It  forms  a  white,  very  volatile  incrustation,  and  has  a 
strong  garlic  odor.     This  indicates  arsenic  or  some  of  its  coin- 
pounds.      (Poison  !) 

(b)  Those  substances  which  form  a  sublimate  in  the  closed 
tube  are  volatile  on  charcoal,  and  some  of  them  give  an  incrus- 
tation which  is  volatile  and  similar  in  color  to  the  sublimate. 

H.   The  substance  burns. 

(a)  The  substance  is  a  metal  which  burns  with  a  brilliant 
white   light,   leaving  a  white   infusible   oxid.      This   indicates 
magnesium. 

(b)  The  substance  is  a  metal  which  burns  with  bright  scintil- 
lations, leaving  a  dark  brown  or  black  oxid,  which  is  magnetic. 
This  indicates  iron. 

(c)  The  substance  is  a  metal  which  burns  with  a  bright  white 
light,  leaving  an  oxid  which  is  yellow  while   hot,  and  white 
when  cold.      This  indicates  zinc. 

(d)  The  substance  burns  with  a  blue  flame,  giving  off  SO2, 
which  may  be  recognized  by  its  odor.     This  indicates  sulfur. 


58  QUALITATIVE  ANALYSIS 

I.  The  substance  is  infusible,  white,  and  highly  luminous  when 
strongly  heated.  It  should  be  allowed  to  cool  somewhat,  and 
the  residue  then  moistened  with  a  drop  or  two  of  a  solution  of 
Co(NO3)2,  and  again  ignited  strongly.  The  mass,  on  cooling, 
should  then  show  a  characteristic  color  as  follows. 

(a)  Slue.     This  indicates  aluminum  oxid  and  alkaline  phos- 
phates or  borates.     Silicon   dioxid  and  certain  silicates  when 
heated  very  strongly  show  this  same  reaction.       * 

(b)  Flesh  color.     This  indicates  magnesium  oxid. 

(<?)  G-reen.  This,  if  yellowish  green,  indicates  zinc  oxid;  if 
bluish  green,  tin  oxid ;  if  a  dirty  dark  green,  antimony  oxid. 

(d)  Violet.     This  indicates  magnesium  arsenate,   borate,  or 
phosphate.     The  latter  is  fusible. 

(e)  Red  brown  while  hot,  colorless  when  cold.     This  indi- 
cates  barium    oxid.     The   residue   gives   an  alkaline  reaction 
with  litmus  paper. 

(/)  G-ray.  This  indicates  an  oxid  of  calcium,  or  strontium, 
the  latter  being  dark  gray.  The  residue  gives  an  alkaline 
reaction  with  litmus  paper. 


ffl.    THE   SUBSTANCE   IS   HEATED    ON   CHARCOAL   WITH   SODIUM 
CARBONATE 

Very  often  the  addition  of  sodium  carbonate  (soda)  assists 
in  the  reduction  of  a  compound,  and  hence  some  of  the  preced- 
ing reactions  are  more  easily  seen  when  this  reagent  is  used. 
The  silicates  are  reduced,  barium  and  strontium  salts  form 
fusible  compounds  which  sink  into  the  charcoal,  while  calcium 
and  magnesium  salts  are  not  changed.  The  salts  of  most  of 
the  other  metals  are  reduced  to  oxids,  or,  in  many  cases,  to 
the  metal  itself,  by  the  combined  action  of  the  soda  and  the 
charcoal. 


REACTIONS  FOR  DRY  SUBSTANCES  59 

(a)  If  a  compound  containing  sulfur  in  any  form  is  strongly 
heated  with  sodium  carbonate  on  charcoal,  it  is  reduced  to 
sodium  sulfid,  Na2S.  If  a  portion  of  the  fused  mass  is  placed 
on  a  clean  silver  coin  and  moistened  with  water,  a  dark  brown 
or  black  stain  of  silver  sulfid  is  produced.  The  particular  way 
in  which  the  sulfur  is  combined  in  the  compound  is  not  indicated 
by  this  action  but  may  be  determined  by  the  action  of  sulfuric 
acid  on  the  compound,  and  by  special  tests.  [See  VIII,  (d).] 

IV.    COLORATION   OF    THE   FLAME 

Many  substances,  especially  the  alkali  and  alkaline  earth  salts, 
impart  a  characteristic  color  to  the  non-luminous  flame.  In 
the  case  of  the  alkali  salts  any  compound  may  be  used.  The 
chlorids  being  most  easily  volatilized  give  the  best  results,  and 
so,  in  case  of  other  salts,  it  is  often  better  to  moisten  the 
compound  with  HC1. 

In  the  examination  a  platinum  wire  is  used.  It  should  be 
cleaned  by  ordinary  methods,  and  then  heated  until  no  color  is 
imparted  to  the  flame.  This  may  often  be  hastened  by  first 
dipping  the  wire  in  HC1.  If  the  wire  has  been  used  with 
some  difficultly  volatile  salt,  it  may  be  necessary  to  heat  in  a 
blast  flame. 

The  sodium  flame  is  nearly  always  visible,  and  very  persist- 
ent, and  not  infrequently  conceals  any  other  color  that  may 
be  present.  Thus  if  sodium  and  potassium  compounds  occur 
together,  the  color  due  to  potassium  cannot  be  seen.  TJie  yel- 
low color  due  to  sodium  compounds  is  entirely  absorbed  if 
the  flame  is  observed  through  a  piece  of  blue  glass.  The  colors 
of  the  other  compounds  may  be  somewhat  modified  by  this,  but 
can  all  be  seen.  For  very  exact  work  a  spectroscope  should  be 
used.  This  shows  the  color  as  a  combination  of  colored  lines, 
which  occupy  a  certain  position  on  a  scale,  and  thus  determine 
with  certainty  the  presence  or  absence  of  any  element.  For 


60  QUALITATIVE   ANALYSIS 

the  scope  of  the  present  work,  however,  this  is  not  necessary. 
The  following  are  the  most  characteristic  colors  produced. 

(a)  A  yellow  flame.     This  indicates  sodium.     The  color  dis- 
appears when  observed  through  blue  glass. 

(b)  A  violet  flame.     This  indicates  potassium,  and  the  color 
is  nearly  the  same  when  observed  through  blue  glass. 

(c)  A    red  flame.      This    indicates    lithium    if    bright   red ; 
strontium  if  carmine  red;  calcium  if  orange  red. 

(d)  A  green  flame.     This  indicates  barium  or  boric  acid  if 
yellowish  green  (the  latter,  if  in  a  salt,  should  first  be  moistened 
with  H2SO4),  copper  if  emerald  green,  thallium  if  bright  grass- 
green.     Phosphates,  if  warmed  with  a  drop  or  two  of  H2SO4, 
show  a  transient  bluish-green  flame. 

(e)  A  blue  flame.     This  indicates  copper  chlorid,  lead,  anti- 
mony, or  arsenic.     Copper  chlorid  and  lead  give  an  azure  blue, 
the  former,  after  heating  a  short  time,  changing  to  green ;  anti- 
mony a  greenish  blue ;  arsenic  a  light  blue. 

V.    COLORATION    OF    THE    BORAX    OR   MICROCOSMIC    BEAD 

Many  metallic  oxids  dissolve  in  fused  borax  or  microcosmic 
salt,  imparting  a  characteristic  color  to  the  fused  mass.  Com- 
pounds with  sulfur,  arsenic,  or  antimony  must  first  be  roasted 
to  change  them  to  oxids.  This  may  be  done  on  charcoal. 
Most  other  salts  are  changed  to  oxids  by  heat,  and  so  show 
this  action. 

A  platinum  wire  is  used,  the  operation  being  as  follows. 
Make  a  small  loop  at  the  end  (not  more  than  two  millimeters 
in  diameter),  heat  to  redness,  and,  while  hot,  plunge  it  into  the 
borax  or  microcosmic  salt,  as  the  case  may  be.  Usually  enough 
of  the  salt  will  adhere  the  first  time ;  if  not,  after  heating,  the 
action  should  be  repeated.  Heat  the  salt  in  the  non-luminous 
flame  until  the  water  of  crystallization  is  driven  off  and  the 


REACTIONS  FOR  DRY  SUBSTANCES  61 

whole  forms  a  clear,  colorless  bead  in  the  loop.  Bring  the 
bead  into  contact  with  the  substance  for  examination,  taking 
care  that  only  a  very  small  quantity  adheres  to  it,  heat  strongly 
in  the  oxidizing  flame,  and  observe  the  color.  It  should  after- 
wards be  heated  in  the  reducing  flame,  and  the  color  observed. 
Either  salt  may  be  used,  although  the  borax  will  usually  be 
found  to  be  the  more  convenient,  the  microcosmic  salt  having  a 
tendency  to  fall  off  from  the  wire.  In  a  few  cases,  however, 
the  latter  must  be  used.  The  difference  in  color  is  very  slight. 
The  color  of  the  bead  while  hot  is  sometimes  characteristic.  The 
following  are  the  most  important  oxids  which  give  colored  beads. 

(a)  Iron  oxid.     In  the  oxidizing  flame  the  bead  is  yellow  or 
deep  red  when  hot,  and  colorless  to  yellow  when  cold.     In  the 
reducing  flame  it  is  green  when  hot,  and  bottle-green  when  cold. 

(b)  Chromium    oxid.     In    the    oxidizing   flame   the    bead   is 
yellowish  green  when  hot,  and  grass-green  when  cold.     In  the 
reducing  flame  it  is  a  fine  emerald  green. 

(<?)  Nickel  oxid.  In  the  oxidizing  flame  the  bead  is  violet 
red  when  hot,  and  brown  when  cold.  In  the  reducing  flame 
it  is  gray  or  opaque  from  the  reduced  nickel,  but  by  long 
heating  it  becomes  colorless. 

(d)  Cobalt  oxid.     In  both  the  oxidizing  and  reducing  flame 
the  bead  is  bright  blue. 

(e)  Manganese  oxid.      In  the  oxidizing  flame  the  bead  is  red- 
dish violet  (amethyst  color).    In  the  reducing  flame  it  is  colorless. 

(/)  Copper  oxid.  In  the  oxidizing  flame  the  bead  is  bluish 
green.  In  the  reducing  flame  it  is  brown;  by  long  heating, 
red  brown. 

(g]  Bismuth  oxid.  In  the  oxidizing  flame  the  bead  is  yel- 
low when  hot,  and  colorless  or  opalescent  when  cold.  In  the 
reducing  flame  it  is  gray  or  opaque  from  the  presence  of 
reduced  metal. 


62  QUALITATIVE  ANALYSIS 

(h)  Silicon  dioxid.  This  dissolves  in  borax  and  gives  a 
colorless  bead  with  either  flame.  In  the  microcosmic  salt  it 
gives  an  insoluble,  opaque  mass,  which  floats  about  in  the  bead 
and  is  known  as  the  "  silica  skeleton." 

VI.    THE   SUBSTANCE   IS   FUSED   ON   PLATINUM   FOIL   WITH   SODIUM 
CARBONATE   AND    POTASSIUM   NITRATE 

This  may  be  done  by  mixing  the  substance  with  three  or 
four  times  as  much  sodium  carbonate,  together  with  a  small 
quantity  of  potassium  nitrate  (saltpeter),  and  heating  on  a 
piece  of  platinum  foil  until  the  mixture  is  completely  fused. 

Only  the  salts  of  manganese  and  chromium  are  to  be  tested 
for  in  this  way,  and  they  should  have  been  indicated  with  more 
or  less  certainty  by  some  of  the  preceding  tests.  The  com- 
pounds of  these  metals  will  be  strongly  oxidized  by  the  potas- 
sium nitrate,  and  the  residue  will  have  a  characteristic  color. 

(a)  The  residue  is  green.     This  indicates  a  manganese  com- 
pound which  has  been  oxidized  to  a  manganate.     The  green 
mass  is  decomposed  by  boiling  water,  and  brown  manganese 
oxid  is  precipitated. 

(b)  The  residue  is  yellow.     This  indicates  a  chromium  com- 
pound which  has  been  oxidized  to  a  chromate.     If  the  color  of 
the  fused  mass  should  not  be  very  distinct,  the  presence  of  the 
chromate  may  be  confirmed  by  dissolving  the  mass  in  warm 
water,  adding  acetic  acid  in  excess  to  decompose  the  excess  of 
sodium  carbonate,  and  precipitating  the  chromic  acid  with  lead 
acetate.     [See  Lead  8.] 

VH.  THE  SUBSTANCE  IS  ACTED  UPON  BY  SULFURIC  ACID 

The  action  of  sulfuric  acid  upon  a  salt  is  primarily  to  liberate 
the  acid  from  which  the  salt  was  derived.  It  quite  often  hap- 
pens that  the  acid  thus  liberated  is  unstable  and  decomposes, 
setting  free  some  gas  which  is  easily  recognized  and  which  is 


REACTIONS  FOR  DRY  SUBSTANCES  63 

characteristic.  A  few  salts,  such  as  sulf ates,  phosphates,  arsenates, 
etc.,  are  not  volatile,  and  so  yield  negative  results  when  treated 
in  this  way.  These  will  have  been  detected  by  some  of  the 
previous  actions,  so  that  the  negative  result  will  be  characteristic. 

Whenever  a  metal  is  acted  upon  by  any  acid  the  primary 
result  is  the  liberation  of  hydrogen.  If  the  acting  acid  is  con- 
centrated sulfuric  acid,  and  the  temperature  is  high,  the  hydro- 
gen which  is  formed  will  immediately  reduce  some  of  the  acid, 
forming  sulfur  dioxid,  which  is  recognized  by  its  odor. 

The  concentrated  acid  is  quite  generally  used,  though  the 
dilute  acid  may  sometimes  have  to  be  employed.  In  the  latter 
case  the  results  are  usually  the  same  if  any  action  takes  place. 
The  differences,  and  the  cases  where  it  is  necessary  to  use  the 
dilute  acid,  will  be  noted  as  they  occur.  In  a  few  cases  explo- 
sive gases  may  be  liberated,  and  in  order  to  avoid  accidents  it 
is  best  to  proceed  as  follows. 

Place  in  a  test-tube  a  small  quantity  of  concentrated  sulfuric 
acid  (not  more  than  a  centimeter  in  depth),  and  heat  to  about 
100°.  Do  not  heat  much  if  any  higher  than  this,  or  the  acid 
itself  may  begin  to  decompose  and  give  off  gases  which  will 
obscure  the  desired  result. 

At  this  temperature  any  explosive  gases  which  may  be  gen- 
erated will  be  decomposed  as  fast  as  formed,  and  so  can  never 
accumulate  in  sufficient  quantities  to  become  dangerous.  Add 
the  substance  under  examination  slowly  and  in  small  quantities, 
and  carefully  observe  the  results. 

The  following  are  the  more  important  gases  given  off  by  this 
action.  When  not  otherwise  stated  they  may  be  recognized  as 
given  on  page  51,  or  by  special  tests. 

(a)  Oxygen.     This  indicates  that  the  substance  was  a  peroxid 
or  some  highly  oxidized  salt,  such  as  a  chromate  or  permanganate. 

(b)  Carbon  dioxid.     This  indicates  a  carbonate,  an  oxalate, 
or  some  organic  compound.     The  latter  generally  blackens,  and 


64  QUALITATIVE   ANALYSIS 

gives  empyreumatic  odors.    [Compare  Carbonic  Acid  4.]    Sulfur 
dioxid  is  also  formed  from  the  presence  of  carbon.     [See  (/).] 

(c)  Carbon    monoxid.      This     indicates    a    formate,    oxalate, 
cyanid,   ferrocyanid,   ferricyanid,    or    some    organic    substance. 
Oxalic  acid  gives  both  CO  and  CO2. 

(d)  Hydrochloric  acid.     This  is  recognized  by  its  suffocating 
odor,  and  the  white  fumes  when  a  glass  rod  dipped  in  ammonia 
is  held  in  the  escaping  gas.     It  indicates  a  chlorid. 

(e)  Hydrobromic  and  hydriodic  acids.     These  present  proper- 
ties very  much  like  hydrochloric  acid,  but  in  addition  are  more 
or  less  completely  decomposed  by  the  heat,  giving  bromin  or 
iodin.     They  indicate  a  bromid  or  iodid  respectively. 

(/)  Hydrofluoric  acid.  This  indicates  a  fluorid.  It  is 
liberated  as  a  colorless  fuming  gas,  which  has  a  suffocating 
odor  and  is  very  corrosive.  If  breathed  in  more  than  minute 
quantities,  it  may  prove  dangerous.  This  acid  etches  glass, 
and  so  when  a  fluorid  is  heated  in  a  clean  test-tube  with 
concentrated  H2SO4  the  sides  of  the  tube  will  be  etched. 

If  a  glass  rod  moistened  with  water  is  held  in  the  escaping 
gas,  it  is  at  once  coated  with  gelatinous  silica,  because  of  the 
decomposition  of  the  SiF4  which  is  present. 

(g)  Hydrocyanic  acid.  This  is  recognized  by  its  odor,  which 
is  similar  to  that  of  KCN.  It  is  exceedingly  poisonous  and 
burns  with  a  violet  flame.  It  indicates  a  cyanid  or  a  sulfo- 
cyanate.  The  ferrocyanids  heated  with  dilute  H2SO4  also 
give  HCN. 

(h)  Hydrogen  sulfid.  This  indicates  a  sulfid.  Dilute  sulfuric 
acid  shows  this  reaction  and  gives  even  better  results  than  the 
concentrated  acid,  for  in  the  presence  of  certain  reducing  agents 
concentrated  H2SO4  may  be  reduced  to  H2S,  especially  at  a  high 
temperature.  [See  VIII,  (c).] 


REACTIONS  FOR  DRY  SUBSTANCES  65 

(i)  Hydrogen.  This  gas  is  inflammable,  and  when  mixed  with 
air  is  explosive,  so  that  when  a  flame  is  brought  to  the  mouth 
of  the  test-tube  in  which  it  is  being  generated  a  sharp  explosion 
often  occurs  (not  dangerous)  and  a  flame  runs  down  the  tube. 
This  indicates  a  metal,  and  the  reaction  is  best  observed  with 
dilute  sulfuric  acid;  for  if  the  acid  is  concentrated  the  hydrogen 
may  reduce  it,  forming  sulfur  dioxid. 

(/)  Sulfur  dioxid.  This  indicates  a  sulfite  or  a  thiosulfate, 
the  latter  giving  in  addition  a  yellow  precipitate  of  sulfur.  This 
action  is  also  given  by  dilute  sulfuric  acid.  Certain  elements, 
such  as  lead,  copper,  mercury,  carbon,  etc.,  when  heated  with 
the  concentrated  acid  also  give  sulfur  dioxid.  [See  (i).] 

(k)  Nitrogen  peroxid.  This  indicates  a  nitrate  or  a  nitrite. 
The  nitrites  are  decomposed  by  cold  dilute  sulfuric  acid,  giv- 
ing the  red  fumes  of  nitrogen  peroxid  at  once,  while  the 
nitrates  give  nitric  acid.  [See  (I).] 

(I)  Nitric  acid.  This  indicates  a  nitrate.  The  nitric  acid 
formed  in  this  way  has  much  the  appearance  of  hydrochloric 
acid,  but  may  be  recognized  by  adding  a  small  piece  of  metallic 
copper,  when  red  fumes  of  nitrogen  peroxid  are  given  off.  [See 
Nitric  Acid  1.]  The  latter  are  also  produced  by  boiling,  which 
decomposes  the  nitric  acid. 

(m)  Chlorin.  This  indicates  a  hypochlorite,  from  which  it 
is  also  liberated  by  dilute  sulfuric  acid. 

(n)  Chlorin  peroxid.  This  is  a  heavy  yellowish-green  gas,  which 
smells  something  like  chlorin  and  is  very  explosive.  It  indicates 
a  chlorate.  By  decomposition  it  forms  chlorin  and  oxygen. 

(o)  Bromin.     This  indicates  a  bromid.     [See  (e).] 

(p)  lodin.  This  indicates  some  iodin  compound,  usually  an 
iodid.  [See  (e).] 

(q)  Acetic  acid.  This  may  be  recognized  by  its  pungent 
odor,  which  is  like  that  of  vinegar.  It  indicates  an  acetate. 


66  QUALITATIVE   ANALYSIS 


Vm.    SPECIAL  TESTS 

In  testing  substances  systematically,  as  in  the  preceding 
sections,  we  sometimes  find  two  or  three  compounds  which 
resemble  each  other  so  closely  as  to  make  it  difficult  to  deter- 
mine what  we  have.  In  order  to  distinguish  such  substances 
we  may  subject  them  to  some  special  and  characteristic  test. 
The  following  are  the  most  important  substances  to  be  distin- 
guished in  this  way. 

(a)  Hydrochloric  acid.     This  gas  fumes  in  the  air  and  has  a 
suffocating  odor.     It  is  in  this  respect  quite  like  hydrobromic, 
hydrofluoric,  and  nitric  acids.     If  to  the  mixture  of  warm  sul- 
furic  acid  and  a  chlorid  [see  VII,  (d)]  a  little  MnO2  is  added, 
chlorin  will  be  liberated,  which  may  be  recognized  by  its  yellow 
color  and  its  peculiar  odor. 

(b)  Hydrobromic  acid.     If  a  mixture  of  sulfuric  acid  and  a 
bromid  be  treated  with  MnO2,  bromin  will  be  liberated,  which 
may  be  recognized  by  its  red-brown  color  and  its  odor,  which  is 
similar  to  that  of  chlorin. 

(c)  Hydriodic  acid.     This  is  also  a  fuming  gas  with  a  suf- 
focating odor;  but  when  liberated  from  an  iodid  with  sulfuric 
acid  it  is  more  or  less  decomposed,  and  iodin  set  free.     The 
latter  is  recognized  by  its  violet-colored  vapor.     Since  hydriodic 
acid  is  a  reducing  agent,  if  the  temperature  is  high  and  the  sul- 
furic acid  concentrated,  the  latter  may  be  reduced  and  hydrogen 
sulfid  liberated.     [See  VII,  (h).] 

(d)  Sulfur  compounds.     These  all  give  the  test  for  sulfur  as 
shown  under  III,  (a).     Having  thus  determined  that  the  com- 
pound contains   sulfur  in   some  form,   we  have   to  determine 
further  what  compound  we  have.     The  following  are  the  most 
important   compounds    of   sulfur  and  the  way  to   distinguish 
them.     Treat  them  with  warm  concentrated  sulfuric  acid  as 


REACTIONS  FOR  DRY  SUBSTANCES  67 

under  VII.  In  most  cases  dilute  sulfuric  acid  will  give  the 
same  reaction. 

Sulfids  will  yield  hydrogen  sulfid,  which  may  be  recognized 
by  its  odor. 

Sulfites  will  yield  sulfur  dioxid,  which  may  be  recognized 
by  its  odor. 

Thiosulfates  will  yield  sulfur  dioxid,  together  with  a  yellow 
precipitate  of  sulfur,  which  remains  undissolved  in  the  acid. 

Sulfates  are  entirely  unaffected  by  both  concentrated  and 
dilute  sulfuric  acid,  and  so  may  be  distinguished  by  this  nega- 
tive action. 

If  a  sulfate  is  strongly  heated  in  a  closed  tube  with  a  small 
piece  of  magnesium  wire,  the  latter  becomes  incandescent,  a 
vigorous  action  takes  place,  MgS  is  formed,  and  SO2  is  liber- 
ated. After  the  tube  is  cool,  if  a  little  dilute  acid  is  added  to 
the  contents,  H2S  is  liberated. 

Sulfocyanates  (called  also  sulfocyanids)  give  free  sulfur,  and 
HCN,  which  may  be  recognized  as  under  VII,  (g). 

[The  rare  elements  selenium  and  tellurium  resemble  sulfur 
very  closely  and  give  a  similar  reaction  when  heated  on  char- 
coal with  sodium  carbonate.  Selenium  when  thus  treated  gives 
an  odor  of  rotten  horseradish.  Tellurium  does  not  give  an  odor, 
but  in  the  closed  tube  gives  a  white  sublimate  which  is  fusible 
but  not  volatile.] 

(e)  Boric  acid.  If  a  borate  is  placed  in  a  test-tube  with  con- 
centrated sulfuric  acid,  a  little  alcohol  added,  and  the  whole 
warmed,  ethyl  borate  is  formed,  which  burns  with  a  character- 
istic green  flame.  If  the  amount  is  small  the  flame  may  appear 
yellow  with  a  green  border. 

(/)  Acetic  acid.  If  an  acetate  is  treated  with  sulfuric  acid, 
acetic  acid  is  liberated.  [See  VII,  (<?).]  If  alcohol  is  also 
added  ethyl  acetate  is  formed,  which  may  be  recognized  by  its 
apple-like  odor. 


68  QUALITATIVE   ANALYSIS 

(g)  Phosphoric  acid.  The  phosphates  are  not  easily  decom- 
posed and  so  the  tests  already  given  are  not  altogether  satisfac- 
tory. [See  II,  I,  (a)  and  IV,  (d).]  If  a  phosphate  is  strongly 
heated  in  a  closed  tube  with  a  small  piece  of  magnesium  wire, 
the  latter  becomes  incandescent,  a  vigorous  action  takes  place, 
and  magnesium  phosphid  is  formed.  After  the  tube  is  cool,  if 
the  contents  are  moistened  with  a  drop  or  two  of  water,  phos- 
phin  is  liberated,  which  has  a  characteristic  and  disagreeable 
odor  something  like  rotten  fish.  If  the  amount  of  phosphin 
is  sufficient,  it  may  take  fire  spontaneously.  Hypophosphites 
when  heated  alone  in  a  closed  tube  will  give  phosphin. 

(h)  Arsenic.  Most  arsenic  compounds  when  heated  on 
charcoal  give  the  characteristic  garlic  odor.  [See  II,  G,  (a).] 
Arsenic  compounds  which  do  not  give  this  when  heated  alone 
should  be  mixed  with  potassium  oxalate  or  potassium  cyanid 
and  heated,  when  the  odor  is  given. 

Most  compounds  of  arsenic  heated  in  a  closed  tube  with  a 
little  charcoal  give  the  characteristic  arsenic  mirror.  [See  I,  C, 
2,  (a).]  All  compounds  of  arsenic  will  give  this  if  previously 
mixed  with  equal  parts  of  sodium  carbonate  and  potassium 
cyanid. 

(i)  Peroxids.  In  addition  to  the  reactions  already  given 
[see  I,  B,  (a)  and  VII,  (a)],  peroxids  and  all  highly  oxidized 
compounds  when  heated  with  concentrated  hydrochloric  acid 
give  free  chlorin. 


PART  III 


SYSTEMATIC    EXAMINATION    FOR    METALS    IN 
.  SOLUTION 

SIMPLE  COMPOUNDS 

This  section  of  the  work  has  for  its  aim  instruction  in  the 
general  scheme  for  qualitative  analysis,  and  is  given  here  because 
it  is  advisable  that  the  analysis  of  simple  compounds  should 
precede  that  of  complex  bodies. 

By  simple  compounds  we  mean  those  which  contain  but  one 
metal  and  one  acid  radical,  the  salts  being  usually  normal.  In 
actual  analysis  there  are  few  simple  compounds,  and  often  the 
substance  to  be  analyzed  is  quite  complex,  containing  several 
metals  and  more  than  one  acid  radical. 

In  the  systematic  analysis  of  substances,  certain  reagents  are 
used  to  separate  the  metals  into  groups,  and  are  therefore  known 
as  group  reagents.  The  groups  are  then  subdivided  by  other 
reagents  until  a  single  metal  is  separated  from  the  others.  The 
separated  metal  is  then  tested  by  methods  given  under  Part  I. 

This  particular  section,  which  treats  of  simple  compounds,  is 
therefore  valuable  in  learning  the  groups  and  group  reagents 
and  how  to  use  them,  as  well  as  the  methods  for  further  sub- 
dividing the  groups.  The  next  section  will  treat  of  the  analysis 
of  mixed  compounds. 

Group  i 

The  solution  should  first  be  tested  with  litmus  paper  to  see 
whether  it  is  neutral,  acid,  or  alkaline.  If  it  is  found  to  be 
alkaline,  it  must  first  be  rendered  acid  by  hydrochloric  acid.  If 

69 


70  QUALITATIVE   ANALYSIS 

it  is  neutral  or  acid,  add  two  or  three  drops  of  HCl  and  observe 
whether  or  not  a  precipitate  is  formed.  If  no  precipitate  forms, 
silver,  mercurous,  and  probably  lead  compounds,  which  consti- 
tute Group  1,  are  absent,  in  which  case  pass  on  to  Group  2.  If 
a  precipitate  forms,  add  more  HCl  until  the  metal  has  been 
completely  precipitated. 

[Certain  basic  salts  of  antimony  or  bismuth  may  be  precipi- 
tated by  the  first  drops  of  HCl.  These  dissolve  easily  on  the 
addition  of  more  acid  and  will  then  appear  in  their  proper  place 
in  the  next  group.] 

Take  about  one-half  of  the  white  precipitate,  add  a  little 
water,  and  boil.  If  the  precipitate  dissolves,  it  is  lead  chlorid. 
Confirm  by  adding  K2CrO4  or  H2SO4. 

[These  and  all  the  following  confirmatory  tests  are  given 
under  the  respective  metals  in  Part  I,  and  should  be  referred  to 
by  the  student  in  testing.] 

To  the  other  half  add  NH4OH.  If  the  precipitate  dissolves, 
it  is  silver  chlorid.  Confirm  by  reprecipitating  with  HNO3. 

If  on  the  addition  of  the  NH4OH  the  precipitate  becomes 
black,  it  is  mercurous  chlorid.  Confirm  by  testing  the  original 
solution  with  SnCL^ 

Group  2 

If  no  precipitate  was  formed  by  the  addition  of  HCl,  add  to 
this  acid  solution  hydrogen  sulfid  until,  after  stirring  or  shaking, 
the  solution  distinctly  smells  of  this  reagent.  (If  the  liquid 
becomes  only  slightly  turbid  by  the  addition  of  the  H2S,  heat 
nearly  to  boiling,  add  more  H2S,  and  allow  it  to  stand  for  a 
short  time.) 

If  no  precipitate  is  formed,  pass  on  to  Group  3.  Group  2 
consists  of  lead,  mercuric,  bismuth,  copper,  cadmium,  arsenic, 
antimony,  and  tin  compounds,  and  the  precipitate,  which  is  a 
sulfid  of  one  of  these  metals,  may  be  yellow,  orange  red,  brown, 
or  black.  If  the  solution  should  be  strongly  acid,  or  contain 


EXAMINATION   FOR  METALS   IN  SOLUTION  71 

some  easily  reducible  compound,  a  milky  appearance,  due  to  the 
separation  of  free  sulfur,  may  be  produced  by  the  reagent. 

(a)  The  precipitate  is  yellow.     Take  a  portion  of  the  precipi- 
tate, add  equal  volumes  of  NH4OH  and  (NH4)2S,  and  warm.    If 
the  precipitate  does  not  dissolve,  it  is  cadmium  sulfid.    Con- 
firm by  testing  the  original  solution  with  NaOH. 

If  the  yellow  precipitate  dissolves  in  (NH4)2S,  it  is  arsenic  or 
stannic  sulfid.  Test  the  original  solution  with  NaOH.  If  no 
precipitate  is  formed,  it  is  arsenic  ;  if  a  precipitate  forms  which 
is  soluble  in  excess  of  the  reagent,  it  is  stannic. 

(b)  The   precipitate  is    orange  red.     It   is    antimony    sulfid. 
Confirm  by  diluting  the  original  solution  with  water,  when  a 
white  basic  compound  will  be  precipitated. 

(c)  The  precipitate  is  brown.     It  is  stannous  or  bismuth  sulfid. 
Take  a  portion  of   the  precipitate  and  treat  it  with  a  little 
(NH4)2SX.     If  it  dissolves,  it  is  stannous.     Confirm  by  testing 
the  original  solution  with  HgCl2. 

If  the  precipitate  is  insoluble  in  (NH4)2SX,  it  is  bismuth.  Con- 
firm by  testing  the  original  solution  with  water  or  with  K2CrO4. 
As  bismuth  sulfid  is  dark  brown  it  may  be  mistaken  for  black, 
but  the  confirmatory  test  will  reveal  its  identity. 

(d)  The  precipitate  is  black.     It  is  lead,  mercuric,  or  copper 
sulfid.     Take  a  portion  of  the  precipitate,  add  an  equal  volume 
of  water  and  a  few  drops  of  concentrated  HNO3,  and  boil  for  a 
moment.     If  it  remains  undissolved,  it  is  mercuric.     Confirm  by 
testing  the  original  solution  with  KI  or  SnCl2. 

If  it  dissolves,  it  is  lead  or  copper.  Add  to  this  solution 
NH4OH  in  excess.  A  light  blue  precipitate,  easily  soluble  in 
the  NH4OH  to  a  deep  blue  solution,  indicates  copper.  This 
needs  no  further  confirmation. 

If  the  addition  of  the  NH4OH  above  produces  a  white  pre- 
cipitate, insoluble  in  excess  of  the  reagent,  it  indicates  lead. 
This  will  have  been  shown  under  Group  1,  unless  the  solution 


72  QUALITATIVE   ANALYSIS 

is  very  dilute.     Confirm  by  testing  the  original  solution  with 
KI  or  K2Cr04. 

[Gold  and  platinum,  which  are  precipitated  by  H2S  as  black 
sulfids,  as  well  as  the  rare  elements  in  other  groups,  are  not 
usually  found  in  an  elementary  course,  and  so  are  omitted.] 

Group^  3 

If  no  precipitate  is  formed  by  the  addition  of  H2S,  take  a 
portion  of  the  original  solution,  add  an  equal  volume  of  NH^Cl, 
and  then  NH4QH  until  it  is  distinctly  alkaline.  If  no  precipi- 
tate is  formed,  pass  on  to  Group  4.  If  a  precipitate  is  formed 
it  is  an  hydroxid  of  either  aluminum,  chromium,  or  iron. 

(a)  The  precipitate  is  white.     It  is  aluminum  hydroxid.     Con- 
firm by  testing  the  original  solution  with  NaOH,  which  should 
give  a  white  precipitate,  soluble  in  excess  of  the  reagent. 

(b)  The  precipitate  is  gray  green.     It  is  chromium  hydroxid. 
Confirm  by  dissolving  in  a  few  drops  of  HNO3,  boiling,  and 
adding  NaOH  in  excess,  which  forms  a  gray-green  precipitate. 

(c)  The  precipitate  is  greenish  white  or  dirty  green,  changing 
to  reddish  brown  by  exposure  to  the  air.     It  is  ferrous  hydroxid. 
Confirm  by  dissolving  in  HNO3,  boiling,  and  adding  NH4OH, 
when  red-brown  ferric  hydroxid  is  precipitated. 

(d)  The  precipitate  is  red-brown.     It  is  ferric  hydroxid.     Con- 
firm by  dissolving  the  precipitate  in  a  few  drops  of  HC1  and 
testing  with  K4Fe(CN)6. 

[The  presence  of  non-volatile  organic  substances,  such  as  tar- 
taric  or  citric  acid,  sugar,  etc.,  prevents  the  precipitation  of  this 
group  by  NH4OH.  These  may  be  removed  by  evaporating  to 
dryness  and  igniting.] 

If  the  solution  contains  phosphoric,  oxalic,  boric,  or  hydro- 
fluoric acid,  combined  with  either  barium,  strontium,  calcium, 
or  magnesium,  the  metal,  with  its  corresponding  acid,  will  be 


,. 


EXAMINATION   FOR  METALS   IN   SOLUTION  73 

precipitated  by  the  addition  of  ammonia,  and  appear  as  a  white 
precipitate  in  the  third  group.  These  acids  must  be  removed 
before  the  metal  can  be  determined.  This  is  a  somewhat  com- 
plicated operation  and  will  be  given  under  Group  3  of  mixed 
compounds. 

Group 

If  no  precipitate  is  formed  by  the  addition  of  NH4OH,  add 
to  this  alkaline  solution  a  slight  excess  of  (NH4)2S  and  heat  to 
boiling.  If  no  precipitate  is  formed,  pass  on  to  Group  5.  If 
a  precipitate  is  formed,  it  is  a  sulfid  of  either  nickel,  cobalt, 
manganese,  or  zinc. 

If  hydrogen  sulfid  gas  is  passed  through  the  solution  con- 
taining NH4OH,  it  forms  (NH4)2S,  and  the  metals  of  this  group 
will  be  precipitated  as  above.  This  method  is  to  be  preferred 
in  many  cases,  since  it  insures  having  freshly  prepared  (NH4)2S, 
which  is  likely  to  decompose  on  standing  and  so  become  worth- 
less. 

(a)  The  precipitate  is  black.     It  is   either  nickel  or  cobalt 
sulfid.     Take  a  portion  of  the  original  solution  and  add  NaOH. 
If  the  precipitate  is  apple  green,  it  is  nickel.     Confirm  by  testing 
with  KCN.     If  the  precipitate  is  blue  and  turns  red  by  boiling, 
it  is  cobalt.     Confirm  by  testing  with  KCN  or  with  KNO2. 

(b)  The  precipitate  is  flesh  colored.     It  is  manganese  sulfid. 
Confirm  by  testing  the  original  solution  with  NaOH  or  NH4OH. 

(c)  The  precipitate  is  white.     It  is  zinc  sulfid.     Confirm  by 
adding  NH4OH  to  the  original  solution  until  the  precipitate 
first  formed  dissolves,  and  then  adding  H2S,  when  a  white  pre- 
cipitate is  formed. 

Group  5 

If  no  precipitate  is  formed  by  (NH4)2S,  add  to  a  portion  of  the 
original  solution  an  equal  volume  of  NH4C1,  enough  NH4OH 
to  render  it  alkaline,  and  then  (NH4)2CO3  to  slight  excess. 


74  QUALITATIVE   ANALYSIS 

If  a  precipitate  is  formed,  it  is  a  carbonate  of  'either  bariumt 
strontium,  or  calcium.  The  precipitate  should  be  white. 

Dissolve  the  precipitate  in  acetic  acid.  To  a  portion  of  this 
solution  add  K2CrO4.  If  a  yellow  precipitate  is  formed,  it  is 
barium  chromate.  Confirm  by  testing  the  original  solution,  con- 
siderably diluted  with  water,  with  H2SO4  or  by  the  coloration 
of  the  flame. 

To  another  portion  add  CaSO4.  If  a  white  precipitate  is 
formed,  it  is  strontium  sulfate.  Confirm  by  testing  the  original 
solution,  diluted  with  an  equal  volume  of  water,  with  H2SO4. 
A  precipitate  should  appear  after  standing  a  short  time.  The 
coloration  of  the  flame  may  also  be  used. 

To  another  portion  add  (NH4)2C2O4.  If  a  precipitate  is  formed 
at  once,  it  is  calcium  oxalate.  Confirm  by  testing  the  or-iginal 
solution  with  H2SO4  after  diluting  with  a  considerable  quantity 
of  water.  If  the  solution  is  sufficiently  diluted,  no  precipitate 
will  appear.  The  coloration  of  the  flame  may  also  be  used. 

Group  6 

This  group  includes  all  those  metals  which  are  not  precipitated 
by  the  preceding  group  reagents.  It  consists  of  magnesium  and 
the  alkali  metals,  and  their  presence  must  be  determined  by 
special  tests. 

If  no  precipitate  is  formed  by  (NH4)2CO3,  add  to  this  alkaline 
solution  some  Na2HPO4.  If  a  white  precipitate  is  formed,  it 
is  magnesium-ammonium  phosphate.  Confirm  the  presence  of 
magnesium  in  the  original  solution  by  any  of  the  tests  for  that 
metal. 

To  the  original  solution  add  a  little  concentrated  NaOH,  and 
boil.  The  odor  of  ammonia,  together  with  an  alkaline  reaction 
in  the  steam  passing  off,  indicates  an  ammonium  compound; 

Add  to  the  original  solution  some  acid  sodium  tartrate,  and 
shake  vigorously  for  a  minute.  A  white  crystalline  precipitate 
indicates  a  potassium  compound.  If  the  solution  is  too  dilute, 


EXAMINATION   FOR  METALS  IN  SOLUTION  75 

it  must  be  concentrated  by  boiling,  or  the  precipitate  may  not 
appear.     Confirm  by  the  coloration  of  the  flame. 

Add  to  the  original  solution  some  acid  potassium  pyroanti- 
monate,  and  shake  well.  A  white  precipitate  indicates  sodium. 
If  the  solution  is  too  dilute,  it  must  be  concentrated  by  boiling, 
or  the  precipitate  may  not  appear.  Confirm  by  the  coloration 
of  the  flame. 

EXAMINATION   FOR  ACID  RADICALS 

In  the  examination  for  acid  radicals  in  solution  we  have  no 
simple  method  of  determination  by  successive  elimination  as  in 
the  examination  for  metals.  The  acids  may  be  grouped  together, 
but  the  same  acid  is  generally  found  in  more  than  one  group. 
The  determination  of  the  metal  should  always  precede  that  of 
the  acid  radical,  and  the  presence  of  certain  metals  in  the  solu- 
tion will  often  eliminate  certain  acid  radicals.  Thus  if  barium 
is  present  the  solution  cannot  contain  sulfuric  acid,  or  if  silver 
is  present  in  a  neutral  or  acid  solution,  the  halogen  acids,  with 
the  exception  of  hydrofluoric  acid,  must  be  absent. 

In  the  analysis  of  simple  compounds,  therefore,  we  first  find 
what  metal  is  present.  Knowing  this  we  may  generally  avoid 
the  necessity  of  looking  for  many  of  the  acid  radicals.  Those 
remaining  may  be  determined  by  the  special  reactions  given  in 
Parti. 


SYSTEMATIC   EXAMINATION    FOR  METALS   IN 
SOLUTION 

MIXED  COMPOUNDS 

In  the  analysis  of  simple  compounds  in  solution  the  method 
of  procedure  is  largely  a  matter  of  convenience.  We  now  take 
up  the  analysis  of  solutions  containing  more  than  one  simple 
compound,  and  our  method  of  procedure  must  be  systematic. 
The  precipitation  and  analysis  of  the  different  groups  must  be 
taken  up  in  their  regular  order ;  for  while  the  group  reagents 
will  precipitate  and  so  separate  the  metals  of  any  group  from 
those  of  the  succeeding  groups,  they  will  not  separate  them 
from  the  preceding  groups  but  will  generally  precipitate  more 
or  less  completely  the  metals  of  those  groups. 

It  is  of  course  understood  that  work  of  this  kind  should 
always  be  carried  on  under  the  direction  of  a  competent  teacher, 
who  should  give  personal  instruction  in  the  various  analytical 
processes.  The  few  general  directions  given  here  will,  if  care- 
fully observed,  save  much  time  and  trouble. 

In  beginning  an  analysis  take  a  sufficient  quantity  of  the 
solution  for  the  purpose  (50  cc.  of  the  solution  will  with  proper 
care  be  found  sufficient  for  most  analyses).  Do  not  take  more 
than  is  necessary,  or  the  different  operations  will  require  more 
time,  and  the  results  will  not  only  be  no  better  but  often  not 
so  clear. 

In  using  the  group  reagents  be  sure  that  the  group  is  com- 
pletely precipitated,  but  avoid  using  a  large  excess  of  the  reagent. 

Before  precipitating  any  group  always  test  the  liquid  with 
the  group  reagent  of  the  preceding  group,  to  be  sure  that  the 
latter  has  been  completely  precipitated. 

76 


EXAMINATION  FOR  METALS  IN  SOLUTION  77 

In  using  any  reagent  for  the  further  division  of  the  group,  or 
in  testing  for  a  particular  metal,  always  avoid  using  a  large 
excess,  unless  this  is  expressly  required,  since  a  large  excess 
may  so  modify  the  results  as  to  render  them  untrustworthy. 

Take  care  that  all  the  analytical  processes  —  precipitation, 
filtration,  washing  the  precipitate,  solution,  etc.  —  are  carefully 
performed.  Many  an  analysis  is  spoiled  through  incomplete 
precipitation  or  incomplete  washing  of  the  precipitate.  When 
results  are  not  what  they  should  be  it  is  certain  that  something 
is  wrong.  Try  to  find  the  cause  of  the  trouble  at  once  ;  failing 
in  this,  ask  the  instructor. 

All  the  different  parts  of  the  analysis  should  be  kept  properly 
labeled  so  that  the  analyst  may  know  what  he  is  doing  and  so 
be  sure  of  his  results.  A  preliminary  test  will  often  show  the 
absence  of  some  metal  in  a  group.  This  will  save  time  which 
would  otherwise  be  spent  in  trying  to  separate  the  metal  from 
the  rest  of  the  group. 

Always  bear  in  mind  that  the  analyst,  if  he  is  successful, 
must  be  careful  in  little  things.  Careful  work  will  always  give 
satisfactory  results,  but  careless  work  is  valueless. 

PRELIMINARY  EXAMINATION 

Before  beginning  the  work  of  separation,  test  the  solution  with 
litmus  paper  to  determine  its  condition.  It  should  be  either  neu- 
tral or  acid.  If  this  is  the  case,  pass  on  to  the  precipitation  of 
Group  1.  If  the  solution  is  neutral,  neither  phosphates,  borates, 
nor  oxalates  will  be  found  in  the  precipitate  of  Group  3. 

Occasionally  an  alkaline  solution  may  be  found,  in  which 
case  it  should  be  acidified  with  HC1.  In  doing  this  a  number 
of  substances  held  in  the  alkaline  solution  may  be  precipitated, 
the  following  being  those  most  likely  to  occur. 

If  a  white  precipitate  appears  which  does  not  dissolve  in  an 
excess  of  the  acid,  it  usually  belongs  to  Group  1  and  may  be 
determined  as  given  under  that  group. 


78  QUALITATIVE   ANALYSIS 

If  the  solution  contains  a  soluble  silicate  (only  the  alkaline 
silicates  are  soluble),  and  is  not  too  dilute,  the  addition  of  HC1 
will  cause  a  partial  precipitation  of  silicic  acid.  This  has  none 
of  the  characteristics  of  the  compounds  found  in  Group  1,  and 
is  gelatinous  in  appearance,  by  which  it  may  generally  be  recog- 
nized. To  determine  what  metals  may  be  present  with  silicic 
acid,  the  whole  solution,  with  an  excess  of  HC1,  must  be  evapo- 
rated to  dryness,  the  dried  mass  extracted  with  water  and  a 
little  HC1,  and  the  metals,  which  will  be  found  in  the  solution, 
separated  in  the  regular  way. 

If  a  white  precipitate  appears  which  dissolves  in  excess  of 
the  HC1,  the  metal  belongs  to  Group  2  or  to  some  subsequent 
group  and  will  appear  in  its  proper  place. 

If  a  yellow  or  reddish-colored  precipitate  is  formed,  the  solu- 
tion contained  a  sulfo-salt  of  arsenic,  antimony,  or  tin,  and 
these  may  be  separated  and  determined  as  given  under  Group  2, 
Subdivision  B. 

If  the  solution  contains  an  alkaline  sulfid,  the  addition  of 
HC1  will  liberate  H2S  and  produce  a  yellowish-white,  finely 
divided  precipitate  of  sulfur. 

In  a  few  other  cases  compounds  may  be  found  dissolved  in 
an  alkaline  solution,  which  give  off  some  characteristic  gas  when 
acidified  with  HC1.  The  gas  will  distinguish  the  acid  in  the 
compound,  the  metal  usually  passing  into  solution  to  reappear 
in  its  proper  place..  In  exceptional  cases  the  precipitate  formed 
may  be  filtered  off,  dried,  and  tested  as  given  under  Blowpipe 

Analysis. 

Group  \ 

To  this  group  belong  Lead,  Silver,  and  Mercury(ous). 

The  metals  are  precipitated  from  a  neutral  or  acid  solution  by 
HC1,  forming  chlorids,  which  are  insoluble  in  dilute  acids. 

The  lead  chlorid,  being  somewhat  soluble,  even  in  cold  water, 
*  is  not  completely  precipitated.  The  part  remaining  in  the  solu- 
tion will  be  precipitated  in  the  second  group. 

^ 


EXAMINATION   FOR  METALS  IN   SOLUTION  79 

If  the  solution  for  any  reason  contains  much  free  nitric  acid, 
any  mercurous  salt  originally  contained  in  the  solution  may  be 
partially  or  entirely  oxidized  to  a  mercuric  compound,  and  so  be 
found  with  the  metals  of  Group  2. 

Take  about  30  cc.  of  the  solution  in  a  beaker  of  convenient 
size,  test  it  with  litmus  paper  and,  if  found  to  be  neutral  or 
acid,  add  to  the  cold  solution  a  few  drops  of  dilute  HC1.  (For 
treatment  of  the  solution  if  found  alkaline,  see  Preliminary 
Examination,  page  77.)  If  no  precipitate  appears,  pass  on  to 
Group  2. 

If  a  precipitate  is  formed,  add  slowly  2  or  3  cc.  of  the  reagent, 
stirring  the  whole  constantly  with  a  glass  rod.  As  the  amount 
of  the  reagent  which  it  is  necessary  to  use  depends  upon  the 
relative  strength  of  reagent  and  solution,  and  also  upon  the 
number  of  metals  to  be  precipitated,  add  the  reagent  as  directed 
and  then  allow  the  precipitate  to  subside.  After  the  supernatant 
liquid  becomes  clear,  add  two  or  three  drops  of  the  reagent, 
allowing  them  to  run  down  the  inside  of  the  beaker,  and  notice 
if  the  liquid  becomes  cloudy.  If  it  does,  the  operation  must  be 
repeated  as  above  until,  after  testing,  the  supernatant  liquid 
remains  clear,  showing  that  the  precipitation  is  complete. 

The  precipitate,  which  should  be  white,  consists  of  the 
chlorids  of  some  or  all  of  the  metals  of  the  group,  and  should 
next  be  filtered  to  separate  it  from  the  surrounding  liquid 
which  contains  the  metals  of  the  remaining  groups.  The  liquid 
which  passes  through  the  filter  is  called  the  filtrate.  This  is 
set  aside  for  later  examination,  and  the  precipitate  examined 
for  Group  1. 

After  transferring  the  precipitate  to  the  filter,  wash  it  twice 
with  cold  water.  This  wash  water  may  be  thrown  away. 

Heat  about  10  cc.  of  water  to  boiling,  and  pour  upon  the  pre- 
cipitate on  the  filter.  If  lead  chlorid  is  present,  it  will  dissolve 
in  the  hot  water  and  pass  through  the  filter.  This  last  filtrate 
may  be  tested  for  lead  by  K2CrO4  or  H2SO4.  [See  Part  I,  Lead.] 


80  QUALITATIVE  ANALYSIS 

If  lead  is  found,  wash  the  precipitate  several  times  with  hot 
water  to  remove  it,  throwing  away  the  wash  water,  and  then 
add  to  the  precipitate  on  the  filter  some  NH4OH.  If  silver 
chlorid  is  present,  it  will  dissolve  in  the  ammonia  and  pass 
through  the  filter.  By  adding  HNO3  in  excess  to  this  alkaline 
filtrate  the  white  silver  chlorid  is  reprecipitated. 

If  mercurous  chlorid  is  present,  it  is  changed  by  the  ammonia 
into  a  black  compound  which  is  a  mixture  of  amido-mercuric 
chlorid,  HgNH2Cl,  and  finely  divided  mercury.  This  is  gener- 
ally regarded  as  sufficient  proof  of  the  presence  of  mercurous 
chlorid,  but  a  further  confirmation  may  be  obtained  by  dissolv- 
ing the  precipitate  in  a  few  drops  of  aqua  regia,  and  testing  the 
mercuric  chlorid  thus  formed  with  SnCl2. 

Group  2 

To  this  group  belong  (Lead),  Mercury (ic),  Bismuth,  Copper,  Cad- 
mium, Arsenic,  Antimony,  and  Tin. 

The  metals  are  precipitated  from  an  acid  solution  by  hydro- 
gen  sulfid,  forming  sulfids,  which  are  insoluble  in  dilute  acids. 
These  sulfids  are  of  two  kinds,  as  shown  by  their  action  when 
treated  with  ammonium  sulfid.  These  form  the  two  subdivisions 
of  the  group. 

Subdivision  A  includes  the  sulfids  of  (Lead),  Mercury(ic), 
Bismuth,  Copper,  and  Cadmium.  These  compounds,  which 
are  basic  in  character,  and  so  sometimes  called  sulfo-bases,  are 
insoluble  in  ammonium  sulfid. 

Subdivision  B  includes  the  sulfids  of  Arsenic,  Antimony,  and 
Tin.  These  compounds,  which  are  acid  in  character,  and  so 
sometimes  called  sulfo-acids,  are  soluble  in  ammonium  sulfid, 
the  latter  compound,  which  is  strongly  basic  in  character, 
uniting  with  these  sulfo-acids  to  form  soluble  sulfo-salts. 

Inasmuch  as  the  analysis  of  this  group  is  attended  with  con- 
siderable difficulty,  it  is  best  to  make  a  preliminary  test  of  the 
solution  to  see  if  any  members  of  the  group  are  present. 


EXAMINATION  FOR  METALS  IN  SOLUTION  81 

Take  a  small  quantity  of  the  acid  filtrate  from  Group  1,  or  of 
the  acid  solution  from  which  Group  1  is  absent,  place  it  in  a 
test-tube  and  treat  it  with  hydrogen  sulfid.  If  no  precipitate 
is  formed,  or  if  only  a  milky,  white  precipitate  is  formed  (which 
is  free  sulfur,  and  caused  by  an  excess  of  acid  in  the  solution), 
the  contents  of  the  test-tube  may  be  thrown  away  and  the 
remainder  of  the  solution  treated  as  if  it  7?ere  the  filtrate  from 
Group  2. 

If  the  solution  contains  nitrates  or  free  nitric  acid,  the  H2S 
will  be  partially  oxidized  to  H2SO4,  and  members  of  Group.  5, 
if  present,  will  be  precipitated.  The  presence  of  Group  5  may 
be  shown  by  adding  a  drop  or  two  of  H2SO4  to  a  few  drops  of 
the  solution  in  a  test-tube.  If  a  precipitate  is  formed,  the  above 
action  may  be  partially  obviated  by  adding  Na2CO3  to  the  nitrate 
from  Group  1  until  a  precipitate  begins  to  form.  Dissolve  this 
precipitate  in  HC1  and  then  proceed. 

If  a  precipitate  was  formed  in  the  test-tube,  the  whole  of  the 
solution  should  be  warmed  to  about  70°  and  a  slow  stream  of 
hydrogen  sulfid  gas  (not  more  than  two  hundred  bubbles  per 
minute)  passed  through  the  warm  liquid.  If  the  solution  has 
not  become  sufficiently  diluted,  enough  warm  water  should  be 
added  to  make  the  total  volume  100  to  150  cc.  before  the  pre- 
cipitation. The  treatment  with  H2S  should  be  continued  until, 
after  blowing  off  the  excess  of  the  gas  from  the  surface  of  the 
liquid,  the  latter  distinctly  smells  of  the  reagent.  This  will 
usually  take  about  ten  minutes. 

Allow  the  precipitate  to  settle.  If  it  does  not  do  so  readily, 
boil  the  whole  for  a  moment.  Decant  the  clear  solution  through 
a  filter,  and  set  aside  this  group  filtrate  for  the  following  groups. 
The  precipitate  is  again  treated  with  boiling  water,  and  after  it 
has  settled  the  whole  is  filtered  and  washed  with  hot  water,  the 
wash  water  being  thrown  away. 

Next  make  another  preliminary  test  to  se.e  if  both  subdivi- 
sions of  the  group  are  present.  This  is  done  by  placing  a  small 


82  QUALITATIVE  ANALYSIS 

portion  of  the  precipitate  in  a  test-tube,  adding  a  few  drops  of 
(NH4)2SX,  and  gently  warming.  If  the  precipitate  is  all  dis- 
solved, only  Subdivision  B  can  be  present.  If  some  of  the  pre- 
cipitate remains  undissolved,  it  belongs  to  Subdivision  A,  in 
which  case,  in  order  to  tell  whether  anything  has  been  dissolved, 
filter,  dilute  with  a  little  water,  and  add  to  the  filtrate  a  slight 
excess  of  dilute  HC1.  A  white,  finely  divided  precipitate  is 
only  sulfur  and  shows  that  Subdivision  B  is  absent,  while  its 
presence  is  shown  by  a  more  or  less  yellow  or  orange-colored 
flocculent  precipitate. 

If  both  subdivisions  are  found  to  be  present,  place  the  whole 
precipitate  in  a  porcelain  dish,  add  enough  (NH4)2SX  to  cover 
the  precipitate,  warm  gently  for  ten  minutes  (do  not  let  the 
mixture  boil),  and  filter.  Wash  two  or  three  times  with  hot 
water,  adding  each  time  a  few  drops  of  the  yellow  ammonium 
sulfid  to  the  precipitate  on  the  filter.  The  precipitate  now 
contains  Subdivision  A,  and  the  filtrate  Subdivision  B. 

Group  2,  Subdivision  A 

Transfer  the  precipitate  containing  Subdivision  A  to  a  porce- 
lain dish  or  a  beaker.  Add  enough  concentrated  HNO3  diluted 
with  an  equal  volume  of  water  to  cover  the  precipitate,  and  boil 
as  long  as  anything  dissolves. 

The  sulfids  of  lead,  bismuth,  copper,  and  cadmium  are  dis- 
solved in  the  nitric  acid,  while  the  mercuric  sulfid,  being  insol- 
uble in  HNO3,  remains  as  a  heavy  black  precipitate.  If  the 
acid  used  is  too  concentrated  and  the  boiling  is  continued  for 
some  time,  the  black  precipitate  may  be  changed  to  a  white 
basic  compound.  [See  Part  I,  Mercurous  4.]  In  the  decomposi- 
tion of  the  sulfids  some  sulfur  is  set  free.  This  appears  as  a 
more  or  less  dark-colored  mass  floating  on  the  surface  of.  the 
liquid.  If  lead  is  present  in  the  original  solution,  most  of  it 
will  have  been  removed  in  Group  1.  A  portion  of  that  remain- 
ing, which  was  precipitated  by  H2S,  may  be  oxidized  by  the 


EXAMINATION  FOR  METALS  IN  SOLUTION  83 

HNO3,'  and  be  found  as  PbSO4  mixed  with  the  insoluble 
mercuric  sulfid,  the  remainder  being  dissolved. 

The  black  precipitate  of  HgS,  or  the  white  basic  compound 
as  noted  above,  may  be  dissolved  on  the  filter  with  a  few  drops 
of  warm  aqua  regia.  Dilute  this  acid  solution  with  water  and 
confirm  the  presence  of  mercury  by  SnCl2  or  KI. 

The  acid  filtrate  from  the  HgS,  which  may  contain  nitrates 
of  lead,  bismuth,  copper,  and  cadmium,  should  be  evaporated 
nearly  to  dryness  to  remove  the  excess  of  HNO3  and  then 
dissolved  in  water.  If  in  dissolving  in  water  a  precipitate 
appears,  add  HNO3  drop  by  drop  until  the  precipitate  (probably 
BiONO3)  disappears.  If  lead  was  found  in  the  first  group,  add 
a  little  (1  cc.)  dilute  H2SO4  and  filter  off  the  precipitated  PbSO4. 

To  this  last  filtrate,  or  to  the  previous  one  if  lead  was  not 
present,  add  an  excess  of  NH4OH.  Bismuth  is  precipitated  as 
white  bismuth  hydroxid,  Bi(OH)3,  while  copper  and  cadmium, 
which  are  first  precipitated,  are  dissolved  in  excess  of  the  reagent. 

The  precipitate  containing  the  bismuth  hydroxid  should  be 
washed  once  or  twice  with  water.  To  confirm  the  presence 
of  bismuth  add  to  the  precipitate  on  the  filter  a  few  drops  of 
dilute  HC1,  and  allow  the  acid  solution  to  filter  into  a  large 
test-tube  filled  with  cold  water.  A  white  precipitate  of  bis- 
muth oxychlorid,  BiOCl,  confirms  the  presence  of  bismuth. 

If  the  alkaline  filtrate,  after  precipitating  the  bismuth,  is  blue, 
it  conclusively  proves  the.  presence  of  copper,  and  may  contain 
cadmium.  If  colorless,  it  can  only  contain  cadmium.  In  the 
latter  case  the  presence  of  cadmium  may  be  confirmed  by  treat- 
ment with  H2S,  when  yellow  cadmium  sulfid,  CdS,  is  precipitated. 

If  the  solution  is  blue,  in  order  to  determine  the  presence  or 
absence  of  cadmium  add  KCN  to  the  solution  until  the  blue 
color  disappears,  and  then  treat  with  H2S,  when  a  yellow  precipi- 
tate proves  the  presence  of  cadmium. 

[It  sometimes  happens  that  the  precipitate  of  CdS  obtained 
in  this  way  is  dark  colored.  This  may  be  due  to  different  causes, 


84  QUALITATIVE   ANALYSIS 

such  as  traces  of  silver  or  mercury  compounds  due  to  incom- 
plete separation  of  the  first  group,  or  failure  completely  to 
remove  members  of  succeeding  groups  by  washing.  In  any 
case  any  considerable  quantity  of  the  precipitate  may  be 
regarded  as  evidence  of  the  presence  of  cadmium.] 

Group  2,  Subdivision  B 

The  alkaline  filtrate  from  Subdivision  A  is  treated  with  dilute 
HC1  until  the  solution  is  acid. 

This  decomposes  the  sulfo-salts,  the  metals  being  precipitated 
as  sulfids,  together  with  more  or  less  free  sulfur.  These  are 
filtered,  washed  with  hot  water,  and  the  filtrate  and  wash  water 
thrown  away. 

It  has  already  been  observed  that  the  metals  which  consti- 
tute Subdivision  B  exist  in  two  states  of  valence  and  so  form 
two  series  of  compounds.  [See  the  respective  metals  in  Part  I.] 
If  compounds  of  these  metals  in  the  lower  form  are  found  in 
the  solution,  the  corresponding  sulfids  are  precipitated  by  H2S. 
These  sulfids  when  dissolved  by  the  yellow  ammonium  sulfid, 
(NH4)2SX,  are  oxidized  (or  sulfurized)  and,  after  decomposition 
of  the  sulfo-salts  by  HC1,  appear  as  the  higher  sulfids.  This  can 
best  be  seen  in  the  case  of  tin,  in  which  the  stannous  sulfid  is 
brown,  while  the  stannic  sulfid  is  yellow. 

The  precipitate  obtained  by  the  decomposition  of  the  sulfo- 
salts  by  HC1,  or  the  whole  group  precipitate  if  Subdivision  A 
was  found  absent  by  the  preliminary  test,  is  placed  in  a  porcelain 
dish  well  covered  with  concentrated  HC1  and  heated  to  boiling. 
(This  last  operation  should  be  carried  on  under  a  hood  on 
account  of  the  large  amount  of  hydrochloric  acid  gas  which 
is  given  off.)  The  whole  should  be  kept  at  or  near  the  boil- 
ing temperature  for  five  minutes  and  filtered  while  hot.  It  is 
better  to  dilute  the  solution  with  a  little  water  before  filtering. 
Arsenic  sulfid  remains  as  a  yellow  precipitate,  while  the  tin  and 
antimony  sulfids  are  dissolved,  forming  chlorids. 


EXAMINATION   FOR  METALS  IN   SOLUTION  85 

Dissolve  the  arsenic  sulfid  by  adding  a  little  concentrated 
HC1  (2  or  3  cc.)  and  a  few  crystals  of  KC1O3.  This  oxidizes 
the  arsenic  compound  to  arsenic  acid.  The  excess  of  HC1 
should  be  removed  by  evaporating  nearly  to  dryness.  Dissolve 
the  nearly  dry  mass  in  a  little  water  (filter  if  necessary),  add  an 
equal  volume  of  NH4C1,  then  NH4OH  to  excess,  and  finally 
MgSO4,  when  white  crystalline  MgNH4AsO4  is  formed.  [See 
Part  I,  Arsenic  Acid  3.] 

The  presence  of  arsenic  may  also  be  shown  by  mixing  a  little 
of  the  yellow  precipitate  with  Na2CO3  and  KCN  and  then  heat- 
ing in  a  closed  tube,  when  the  "  arsenic  mirror  "  will  be  formed. 
[See  Part  II,  VIII,  (h).] 

The  acid  filtrate  from  the  arsenic  sulfid,  which  may  contain 
both  antimony  and  tin  in  the  form  of  chlorids,  must  first  be 
tested  to  see  if  antimony  is  present.     This  can  often  be  deter- 
mined by  the  color  of  the  precipitate  of  Subdivision  B,  since 
the  orange-red   color  of  the  antimony  sulfid  is  characteristic. 
A  small  quantity  may,  however,  be  overlooked,  and  so,  to  make 
sure,  place  a  drop  of  the  acid  solution  on  a  clean  platinum 
foil   and  put  into  it  a  bit  of   metallic  zinc.     If   antimony  is    0 
present,   a  black    spot    of    metallic    antimony  is    immediately    ^ 
formed  on  the  foil.     This  is  soluble  in  HNOo. 

o 

If  antimony  is  found  to  be  present,  place  the  whole  of  the 
acid  solution  in  a  porcelain  dish,  put  into  it  a  piece  of  clean 
platinum  foil  and  a  piece  of  zinc  in  such  a  way  that  the  two 
metals  shall  be  in  contact. 

The  antimony  which  is  present  in  the  solution  as  the  chlorid, 
SbCl3,  will  be  reduced  to  metallic  antimony  and  form  a  black 
deposit  upon  the  platinum  foil. 

The  tin,  which,  if  present,  is  in  the  solution  as  stannic  chlorid, 
SnCl4,  will  be  reduced  to  stannous  chlorid,  SnCl2,  or,  if  the 
action  is  continued  long  enough,  to  metallic  tin.  The  latter 
will  be  found  as  a  gray  deposit  on  and  about  the  zinc.  Filter  sS-K 
and  test  the  filtrate  with  HgCl2.  A  white  precipitate  indicates 
tin.  [See  Part  I,  Mercuric  8.] 


86  QUALITATIVE  ANALYSIS 

If  the  action  is  allowed  to  continue  until  all  the  tin  in  the 
solution  is  reduced  to  the  metallic  state,  the  white  precipitate 
will  not  appear.  If  this  should  be  the  case,  as  a  further  pre- 
caution pour  off  the  solution  as  carefully  as  possible  from  the 
deposited  metals,  remove  any  zinc  which  may  remain  undis- 
solved,  add  to  the  metallic  mixture  a  little  concentrated  HC1, 
and  warm.  If  tin  is  present,  it  will  dissolve,  forming  stannous 
chlorid,  which  'may  be  filtered  from  the  undissolved  antimony 
and  tested  as  above. 

[If  it  is  desirable  to  know  whether  the  tin  was  originally 
present  as  a  stannous  or  a  stannic  compound,  take  a  few  drops 
of  the  original  solution,  remove  any  of  the  metals  of  Group  1 
which  may  have  been  present,  and  add  a  little  HgCl2.  A  white 
or  gray  precipitate,  insoluble  in  concentrated  HC1,  proves  the 
tin  to  be  stannous.] 

Group  3 

To  this  group  belong  Aluminum,  Chromium,  and  Iron. 

If  phosphates,  oxalates,  silicates,  borates,  or  fluorids  are  pres- 
ent in  the  solution,  the  corresponding  salts  of  Barium,  Strontium, 
Calcium,  and  Magnesium  may  be  precipitated  with  this  group. 

The  presence  of  these  salts  very  materially  complicates  the 
whole  treatment  of  the  group,  and  so,  in  making  up  solutions 
for  an  elementary  analytical  course,  they  should  be  omitted 
until  the  student  has  become  familiar  with  the  separations 
without  them.  The  method  of  treatment  when  these  salts  are 
present  will  be  found  on  page  90. 

The  metals  of  this  group  are  precipitated  by  ammonium 
hydroxid,  in  the  presence  of  ammonium  chlorid,  forming 
hydroxids. 

In  addition  to  the  metals  of  this  group  mentioned  above,  if 
manganese  is  present  in  the  solution,  traces  of  it  are  always 
found  in  this  group  precipitate.  This  is  due  to  the  fact  that 
manganese,  while  not  immediately  acted  upon  by  NH4OH  in 


EXAMINATION   FOR  METALS  IN  SOLUTION  87 

the  presence  of  NH4C1,  gradually  becomes  oxidized  and  is  then 
more  or  less  completely  precipitated  by  the  NH4OH,  especially 
after  long  standing.  [See  Part  I,  Manganese  2.] 

Before  precipitating  Group  3  the  nitrate  from  Group  2  must 
be  boiled  until  the  H2S  remaining  in  the  solution  has  been  com- 
pletely expelled.  If  this  is  not  done,  the  NH4OH,  which  is 
used  to  precipitate  Group  3,  will  form  (NH4)2S  with  the  H2S, 
and  any  metals  belonging  to  Group  4  which  may  be  present 
will  be  precipitated  by  it.  The  boiling  must  be  continued  until 
the  odor  of  H2S  entirely  disappears,  or,  better,  until  a  piece  of 
filter  paper  moistened  with  lead  acetate,  when  held  in  the 
escaping  steam,  remains  unchanged,  it  being  blackened  by  H2S. 
If  Group  2  was  not  present,  the  solution  need  only  be  heated 
to  the  boiling  point. 

If  iron  is  present,  it  may  have  been  in  the  original  solution 
either  as  a  ferrous  or  a  ferric  compound,  but  before  precipita- 
tion in  this  group  it  must  always  be  in  the  ferric  state,  since 
most  ferrous  compounds  are  unstable  and  not  completely  pre- 
cipitated by  the  group  reagent.  Furthermore,  whatever  may 
have  been  the  condition  of  the  iron  in  the  original  solution,  if 
it  has  been  acted  upon  by  H2S,  as  it  must  have  been  if  the 
second  group  was  present,  all  of  the  iron  will  now  be  in  the 
ferrous  condition,  because  of  the  reducing  action  of  the  H2S, 
and  so  it  will  always  have  to  be  oxidized  before  precipitation. 
[See  Part  I,  Ferric  3.] 

A  preliminary  test  for  iron  should  be  made  by  taking  a  few 
drops  of  the  solution  in  a  test-tube,  adding  two  or  three  drops 
of  HNO3,  boiling,  and  then  testing  with  K4Fe(CN)6.  A  blue 
precipitate  indicates  iron.  [See  Part  I,  Ferric  7.] 

If  iron  is  found,  add  to  the  hot  solution  a  little  concentrated 
HNO3  (not  more  than  1  or  2  cc.)  and  boil.  If  the  solution 
becomes  dark  colored,  continue  adding  the  HNO3  to  the  hot 
solution  drop  by  drop  until  it  becomes  clear,  when  the  oxidation 
will  be  complete. 


88  QUALITATIVE   ANALYSIS 

[Be  careful  not  to  add  more  HNO3  than  is  necessary,  since 
manganese,  if  present,  will  also  be  oxidized,  and  will  then  be 
precipitated  with  Group  3.] 

Add  to  the  solution  an  equal  volume  of  NH4C1,  to  keep  metals 
belonging  to  Group  4  from  precipitating,  and  then  NH4OH  to 
slight  but  distinct  excess.  Boil  for  a  moment  and  filter  as  soon 
as  the  precipitate  settles.  The  filtrate,  which  contains  the  metals 
of  the  succeeding  groups,  is  set  aside  for  further  examination. 
The  precipitate  contains  the  hydroxids  of  the  metals  of  the  third 
group,  together  with  certain  phosphates,  oxalates,  etc.,  if  these 
salts  were  present  in  the  solution. 

PHOSPHATES,  OXALATES,  ETC.,  ARE  ABSENT 

If  the  color  of  the  precipitate  is  white,  it  can  consist  of  alumi- 
num hydroxid  only.  If  it  is  gray  green,  chromium  hydroxid, 
and  perhaps  aluminum  hydroxid,  is  present.  If  it  is  dark  red- 
brown,  all  of  the  metals  of  the  group  may  be  present. 

Dissolve  the  precipitate  in  warm  dilute  HC1,  avoiding  a  large 
excess  of  the  acid.  Add  an  excess  of  a  concentrated  solution 
of  sodium  hydroxid  and  heat  to  boiling,  keeping  the  whole  at 
that  temperature  for  a  short  time. 

All  the  metals  of  the  group  are  precipitated  as  hydroxids,  but 
the  aluminum  hydroxid  is  dissolved  in  excess  of  the  reagent, 
forming  sodium  aluminate,  NaAlO2.  Filter  and  wash  the  pre- 
cipitate with  hot  water. 

Add  to  the  alkaline  filtrate  HC1  until  it  is  distinctly  acid,  and 
then  add  NH4OH  until  it  is  alkaline.  An  almost  transparent, 
flocculent  precipitate,  which  at  first  usually  rises  to  the  surface 
of  the  liquid,  indicates  the  presence  of  aluminum. 

[If  the  original  solution  contains  tin  in  the  form  of  stannous 
chlorid,  and  too  much  HC1  was  added  before  precipitation  of 
the  second  group,  some  of  the  stannous  sulfid  there  precipitated 
may  have  been  dissolved,  forming  stannous  chlorid.  This  will 
appear  in  the  solution  with  the  aluminum  and  be  precipitated 


EXAMINATION   FOR  METALS   IN   SOLUTION  89 

with  it.  [See  Part  I,  Stannous  2.]  In  order  to  detect  the  tin, 
heat  a  small  portion  of  the  alkaline  nitrate  to  boiling  and  then 
add  a  drop  or  two  of  a  solution  of  Bi(NO3)3.  A  black  precipi- 
tate of  metallic  bismuth  shows  the  presence  of  the  stannous 
chlorid.  [See  Part  I,  Bismuth  8.]  Instead  of  the  alkaline 
filtrate  a  little  of  the  aluminum  hydroxid  precipitate  may  be 
dissolved  in  NaOH  and  then  treated  as  above. 

If  tin  is  found  by  this  test,  in  order  to  determine  whether 
aluminum  is  present,  render  the  whole  alkaline  filtrate  very 
slightly  acid  with  HC1 ;  or  dissolve  the  white  precipitate  in  a 
very  slight  excess  of  HC1,  dilute  to  about  100  to  150  cc.,  and 
precipitate  the  tin  with  H2S.  If  the  solution  is  sufficiently 
dilute,  the  stannous  sulfid  will  not  dissolve  in  the  HC1.  Filter 
off  the  precipitate,  if  any,  and  treat  the  filtrate  as  before  to 
detect  the  presence  of  aluminum.] 

The  precipitate  formed  by  NaOH  may  contain  both  iron  and 
chromium,  and  possibly  traces  of  manganese. 

If  the  precipitate  has  a  dark  red-brown  color,  this  is  usually 
considered  sufficient  evidence  that  iron  is  present.  If  further 
proof  of  its  presence  is  desired,  dissolve  a  small  portion  of 
the  precipitate  in  HC1,  and  test  the  solution  with  K4Fe(CN)6 
or  KSCN. 

The  presence  of  chromium  may  be  shown  by  fusing  a  small 
portion  of  the  precipitate,  mixed  with  sodium  carbonate  and 
potassium  nitrate,  on  a  piece  of  platinum  foil.  This  oxidizes 
the  chromium  to  a  chromate,  which  appears  as  a  yellow  mass 
on  the  foil.  [See  Part  II, VI.] 

If  manganese  is  present,  it  will  also  become  oxidized  to  a 
manganate,  forming  a  green  mass  on  the  foil.  Since  the  green 
manganate  entirely  conceals  the  yellow  chromate,  in  order  to 
show  the  presence  of  chromium  when  manganese  is  also  present, 
dissolve  the  fused  mass  in  warm  water,  filter  if  necessary,  add 
an  excess  of  acetic  acid,  and  boil  to  decompose  the  sodium  car- 
bonate. The  green  manganate  is  more  or  less  completely 


90  QUALITATIVE   ANALYSIS 

decomposed,  while  the  chromate  dissolves  unchanged.  If  lead 
acetate  is  now  added  to  the  acid  solution,  yellow  lead  chro- 
mate, PbCrO4,  is  precipitated,  which  confirms  the  presence  of 
chromium. 

[It  is  often  desirable  to  know  whether  the  iron  is  present  as 
a  ferrous  or  ferric  compound.  To  determine  this,  take  a  little 
of  the  original  solution  and,  if  acid,  nearly  or  quite  neutralize 
it  with  a  concentrated  solution  of  Na2CO3.  If  a  slight  precipi- 
tate forms,  dissolve  it  with  a  drop  or  two  of  HC1.  Add  to  the 
neutral  or  very  slightly  acid  solution  freshly  prepared  barium 
carbonate,  BaCO3  (which  should  be  previously  shaken),  in  suffi- 
cient quantity  to  precipitate  any  aluminum,  chromium,  or  ferric 
iron  which  may  be  present.  The  operation  is  best  carried  on  in 
a  flask,  which  should  be  loosely  corked.  Shake  the  mixture 
occasionally  for  ten  minutes,  allow  the  precipitate  to  settle,  and 
filter.  The  filtrate  now  contains  all  the  ferrous  iron,  and  the 
precipitate  all  the  ferric.] 

PHOSPHATES,  OXALATES,  ETC.,  ARE  PRESENT 

If  the  original  solution  was  neutral  or  alkaline,  none  of  these 
salts  can  be  present.  If  it  was  acid,  any  or  all  of  them  may  be 
present. 

Precipitate  the  group  as  given  above  when  these  salts  are 
absent,  observing  all  the  precautions  there  mentioned.  Filter 
and  set  aside  this,  the  original  group  filtrate,  for  later  examina- 
tion, since,  if  the  amount  of  phosphates,  oxalates,  etc.,  is  small, 
it  may  contain  metals  belonging  to  each  of  the  succeeding 
groups.  Preliminary  tests  should  then  be  made  for  each  acid 
before  making  the  separation. 

Phosphoric  acid  is  best  shown  by  taking  a  small  portion  of 
the  precipitate  in  a  test-tube,  dissolving  in  a  little  concentrated 
HNO3,  adding  ammonium  molybdate,  and  warming.  A  yellow 
precipitate,  insoluble  in  HNO3,  indicates  phosphoric  acid.  [See 
Part  I,  Phosphoric  Acid  5.] 


EXAMINATION   FOR  METALS   IN  SOLUTION  91 

Oxalic  acid  is  best  shown  by  dissolving  a  small  portion  of  the 
precipitate  in  HNO3,  adding  an  excess  of  a  concentrated  solu- 
tion of  Na2CO3,  and  boiling.  The  metals  present  are  all  precipi- 
tated, and  the  oxalic  acid  combines  with  the  reagent  to  form 
sodium  oxalate.  Filter,  add  to  the  filtrate  acetic  acid  in  excess, 
boil  to  expel  the  CO2,  and  add  CaCl2.  A  white  precipitate, 
insoluble  in  acetic  acid,  indicates  oxalic  acid.  [See  Part  I, 
Oxalic  Acid  4.] 

The  three  other  acids  are  not  likely  to  be  met  with  in  an 
elementary  course ;  in  more  advanced  practical  work  they  are 
quite  often  found,  especially  silicic  acid,  which  is  found  in 
nearly  all  minerals. 

Silicic  acid  may  be  shown  by  making  a  microcosmic  bead  and 
introducing  a  bit  of  the  precipitate.  A  "  silica  skeleton  "  indi- 
cates silicic  acid.  [See  Part  II,  V,  (h).] 

Boric  acid  is  best  shown  by  the  flame  test.  Take  a  small 
portion  of  the  precipitate  in  a  porcelain  dish,  add  concentrated 
H2SO4  and  a  little  alcohol,  warm,  and  ignite.  A  green  or  green- 
bordered  flame  indicates  boric  acid.  [See  Part  II,  VIII,  (e).~\ 

Hydrofluoric  acid  is  best  shown  by  the  etching  of  glass. 
[See  Part  II,  VII,  (/).] 

Having  shown  the  presence  of  any  of  these  acids,  before 
separating  the  metals  they  must  be  removed. 

The  whole  group  precipitate  is  dissolved  in  a  slight  excess 
of  HC1  and  then  reprecipitated  by  concentrated  NaOH.  The 
aluminum  precipitate,  whether  present  in  the  form  of  hydroxid 
or  phosphate,  is  dissolved  in  excess  of  the  reagent,  and  its 
presence  confirmed  as  before.  [See  page  88.] 

Silicic  acid  will  rarely  be  found  in  a  solution  intended  for  an 
elementary  analytical  course.  If,  however,  it  should  be  found, 
it  must  be  removed  first  of  all. 

To  remove  silicic  acid,  dissolve  the  original  group  precipitate 
in  HNO3,  and  evaporate  the  whole  to  complete  dryness  on  a 
water  bath.  This  changes  any  silicic  acid  which  may  be  present 


92  QUALITATIVE  ANALYSIS 

into  silicon  dioxid,  SiO2,  which  is  insoluble  in  water  and  in  acids. 
Add  to  the  dried  mass  hot  water  and  a  little  HNO3.  The  metals, 
which  are  now  in  the  form  of  nitrates,  will  dissolve,  while  the 
silicon  dioxid  will  remain  undissolved.  Filter  off  the  insoluble 
silicon  dioxid,  which  may  be  tested  and  then  thrown  away,  but 
keep  the  acid  filtrate.  If  either  phosphoric  or  oxalic  acid  was 
found  by  the  preliminary  tests,  it  must  next  be  removed.  [See 
below.]  If  neither  of  these  acids  was  found  to  be  present,  add 
NH4C1  to  the  acid  nitrate,  precipitate  the  group  with  NH4OH, 
and  separate  the  metals  as  already  given.  The  nitrate  from  this 
last  precipitation  should  be  added  to  the  original  group  nitrate 
for  further  examination. 

To  remove  phosphoric  acid,  take  the  acid  nitrate  after  remov- 
ing the  silicic  acid,  or  if  silicic  acid  was  not  present,  take  the 
NaOH  precipitate,  and  after  dissolving  it  in  a  little  concentrated 
HNO3  in  a  porcelain  dish,  add  a  sufficient  quantity  of  pure  tin 
foil  (a  piece  about  four  inches  square  will  generally  be  sufficient), 
and  heat  to  boiling.  Allow  the  hot  mass  to  stand  a  few  minutes, 
dilute  with  an  equal  volume  of  water,  and  filter.  The  concen- 
trated HNO3  acts  upon  the  tin,  forming  metastannic  acid,  which 
combines  with  the  phosphoric  acid,  forming  a  compound  of  some- 
what doubtful  composition.  This  latter  compound  is  insoluble 
in  HNO3,  and  so  all  the  phosphoric  acid  will  be  found  in  the 
precipitate.  The  nitrate  should  now  be  tested  to  see  if  phos- 
phoric acid  is  still  present.  If  found,  the  operation  must  be 
repeated  until  it  has  been  entirely  removed.  If  oxalic  acid  was 
found  by  the  preliminary  test,  it  must  now  be  removed.  If  no 
oxalic  acid  is  present,  add  NH4C1  to  the  acid  filtrate,  precipitate 
the  group  with  NH4OH,  and  separate  the  metals  as  before.  The 
nitrate  from  this  last  precipitation,  which  may  contain  metals 
belonging  to  the  succeeding  groups,  should  be  added  to  the 
original  group  nitrate  for  further  examination. 

To  remove  oxalic  acid,  take  the  acid  filtrate  after  remov- 
ing the  phosphoric  acid,  or  if  phosphoric  acid  was  not  present, 


EXAMINATION   FOR  METALS  IN  SOLUTION  93 

take  the  NaOH  precipitate,  and  after  dissolving  it  in  a  little 
HNO3  add  a  concentrated  solution  of  Na2CO3  in  slight  excess, 
and  boil  for  a  moment.  The  oxalates  are  all  decomposed  and 
form  soluble  sodium  oxalate,  the  metals  all  being  found  in  the 
precipitate  in  the  form  of  carbonates  or  hydroxids.  Filter  and 
wash  with  hot  water.  The  nitrate  contains  the  sodium  oxalate, 
and  after  it  has  been  tested  may  be  thrown  away.  Dissolve 
the  precipitate  in  HC1,  add  NH4C1,  precipitate  the  group  with 
NH4OH,  and  separate  the  metals  as  before.  The  filtrate  from 
this  last  precipitation  should  be  added  to  the  original  group 
filtrate  for  further  examination. 

Boric  and  hydrofluoric  acids  are  very  difficult  to  separate. 
They  are  found  in  small  quantities  in  a  large  number  of  minerals, 
but  will  hardly  be  found  in  an  elementary  analytical  course. 
The  method  of  removing  them  belongs  therefore  to  a  more 
advanced  course  and  so  will  be  omitted  here. 

Group  4 

To  this  group  belong  Nickel,  Cobalt,  Manganese,  and  Zinc. 

The  metals  are  precipitated  from  a  solution  containing  NH4C1 
and  an  excess  of  NH4OH,  by  means  of  hydrogen  or  ammonium 
sulfid,  forming  sulfids.  These  are  all  insoluble  in  alkalies,  but 
a  part  of  them  are  easily  soluble  in  acids. 

Since  the  precipitation  and  separation  of  the  metals  of  this 
group  is  attended  with  considerable  difficulty,  it  is  always  best 
to  make  a  preliminary  test  to  see  if  any  of  them  are  present. 
This  is  best  done  by  adding  a  few  drops  of  (NH4)2S  to  a  small 
portion  of  the  solution.  If  a  precipitate  forms,  some  of  the 
metals  of  the  group  are  present,  and  the  whole  solution  must  be 
treated  in  a  somewhat  similar  way. 

The  filtrate  from  Group  3  usually  contains  enough  NH4C1. 
Add  NH4OH,  if  necessary,  until  the  liquid  smells  strongly  of 
the  reagent,  and  pass  a  slow  stream  of  hydrogen  sulfid  gas 
through  the  liquid  until  it  is  saturated. 


94  QUALITATIVE   ANALYSIS 

Instead  of  the  H2S,  colorless  ammonium  sulfid,  (NH4)2S,  is 
very  commonly  used  to  precipitate  this  group.  This  reagent 
decomposes  on  standing,  and  often  contains  notable  quantities 
of  the  yellow  ammonium  sulfid,  (NH4)2SX.  For  these  reasons 
hydrogen  sulfid  is  recommended  as  the  better  and  more  con- 
venient reagent  to  use.  This  forms  (NH4)2S  with  the  excess 
of  NH4OH  in  the  solution,  and  so  the  final  result  is  the  same 
whichever  reagent  is  used. 

After  precipitation  the  whole  should  be  boiled  in  a  flask  until 
the  excess  of  (NH4)2S  has  been  decomposed  and  the  liquid  no 
longer  smells  of  the  reagent,  or  at  most  only  slightly.  This 
may  take  some  minutes,  but  should  always  be  done.  The  pre- 
cipitate should  now  settle  quickly,  leaving  a  clear  solution  above. 
[Sometimes  the  liquid  above  the  precipitate  appears  dark  brown 
in  color.  This  is  due  to  the  fact  that  NiS  is  slightly  soluble  in 
(NH4)2S,  giving  a  dark  brown  solution.  [See  Part  I,  Nickel  5.] 
If  this  is  the  case,  or  if  the  precipitate  does  not  settle  quickly, 
continue  the  boiling  for  a  little  time,  adding  water  and  a  little 
NH4OH  to  replace  that  lost  by  evaporation,  until  the  precipitate 
settles  quickly  and  the  solution  becomes  clear.] 

Filter  quickly  while  hot,  and  preserve  the  filtrate  for  further 
examination,  after  testing  it  with  (NH4)2S  to  see  if  the  group 
has  been  completely  precipitated.  The  precipitate  is  washed 
with  hot  water,  and  the  wash  water  thrown  away. 

The  sulfids,  particularly  those  of  nickel  and  cobalt,  have  a 
tendency  to  oxidize  on  the  filter  and  thus  form  sulfates.  The 
latter  are  soluble  in  water  and  so  pass  through  with  the  wash 
water.  If  the  boiling  is  continued  long  enough,  and  in  a  flask, 
and  if  the  filtering  takes  place  while  the  solution  is  still  hot, 
the  sulfids  will  not  usually  oxidize  enough  to  do  any  harm.  If 
oxidation  should  take  place  (and  this  can  usually  be  noted  by 
the  grayish  film  which  appears  on  the  surface  of  the  black 
precipitate),  a  little  H2S  water  or  (NH4)2S  added  to  the  filter 
will  stop  it. 


EXAMINATION   FOR  METALS  IN  SOLUTION  95 

If  the  precipitate  is  white,  it  consists  of  zinc  sulfid  only.  If 
it  is  flesh  colored,  it  contains  manganese  sultid,  and  perhaps  zinc 
suliid.  If  it  is  black,  all  the  metals  of  the  group  may  be  present. 

Make  a  hole  through  the  bottom  of  the  filter  in  the  funnel 
and  wash  the  contents  into  a  beaker.  Add  to  the  sulfids  in  the 
beaker  about  20  to  30  cc.  of  dilute  HC1,  warm  for  five  minutes 
without  boiling,  filter,  and  wash.  The  manganese  and  zinc  sulfids 
are  easily  dissolved  by  the  dilute  acid  and  will  be  found  in  the 
filtrate  as  chlorids,  while  the  nickel  and  cobalt  sulfids,  being 
insoluble  in  the  dilute  acid,  remain  in  the  precipitate.  [If  the 
acid  used  is  too  concentrated,  a  very  little  of  the  cobalt  sulfid 
may  dissolve  and  give  a  faint  pink  color  to  the  solution.  This 
will  do  no  harm.] 

The  black  precipitate,  which  may  contain  both  nickel  and 
cobalt  sulfids,  is  next  examined  with  the  borax  bead  to  see  if 
cobalt  is  present.  A  blue  bead  confirms  the  presence  of  cobalt. 
Dissolve  the  precipitate  in  a  small  quantity  of  aqua  regia, 
evaporate  nearly  to  dryness,  and  dissolve  the  nearly  dry  mass 
in  water.  If  no  cobalt  was  found  by  the  borax  bead,  the  nickel 
is  now  precipitated  from  the  solution  with  NaOH,  forming 
apple-green  Ni(OH)2.  If  cobalt  was  found  by  the  borax  bead, 
take  a  portion  of  the  solution,  add  KCN  until  the  precipitate 
first  formed  dissolves  in  excess,  and  boil  for  a  minute.  Next 
add  a  little  NaOH,  and  finally  bromin  water  in  excess.  A  black 
precipitate  indicates  nickel,  since  cobalt  does  not  give  this  reac- 
tion. [See  Part  I,  Nickel  7  and  Cobalt  7.] 

These  preliminary  tests  for  the  presence  of  nickel  and 
are  usually  sufficient  for  a  qualitative  analysis. 

If  it  is  desirable  to  separate  the  two,  add  to  the  aqueous  solu- 
tion, after  having  dissolved  the  sulfids  in  aqua  regia,  acetic  acid, 
and  then  potassium  nitrite  in  excess.  The  whole  should  be 
warmed  and  allowed  to  stand  several  hours.  The  cobalt  is 
completely  precipitated  as  yellow  Co(NO2)3(KNO2)3,  the  nickel 
remaining  in  solution.  [See  Part  I,  Nickel  10  and  Cobalt  10.] 


96  QUALITATIVE  ANALYSIS 

Filter  and  precipitate  the  nickel  from  the  filtrate  with  NaOH. 
Confirm  the  presence  of  nickel  in  the  last  precipitate  by  means 
of  the  borax  bead.  [See  Part  II,  V,  (e).] 

The  acid  filtrate,  containing  the  manganese  and  zinc  in  the 
form  of  chlorids,  should  be  boiled,  if  necessary,  until  no  more 
H2S  is  given  off,  and  then  treated  with  a  slight  excess  of  con- 
centrated NaOH.  Both  metals  are  precipitated  as  hydroxids, 
but  the  zinc  hydroxid  dissolves  in  excess  of  the  reagent.  Filter 
and  examine  the  filtrate  with  H2S.  A  white  precipitate  of  ZnS 
confirms  the  presence  of  zinc. 

The  precipitate  of  manganese  hydroxid,  which  should  at  first 
be  white,  may  be  dark  colored  from  a  trace  of  cobalt  dissolved 
by  the  HC1.  In  any  case  it  becomes  dark  brown  after  standing 
a  few  minutes.  This  is  caused  by  oxidization  and  the  forma- 
tion of  manganic  hydroxid,  Mn(OH)3,  which  is  usually  con- 
sidered sufficient  proof  of  the  presence  of  manganese.  [See 
Part  I,  Manganese  1.]  A  further  confirmatory  proof  of  the 
presence  of  manganese  may  be  obtained  by  fusing  a  portion  of 
the  precipitate  on  a  piece  of  platinum  foil  with  Na2CO3  and 
KNO3,  forming  a  green  manganate. 

Group  5 

To  this  group  belong  Barium,  Strontium,  and  Calcium. 

The  metals  are  precipitated  from  a  solution  containing  an 
excess  of  NH4OH  by  means  of  ammonium  carbonate,  forming 
carbonates. 

filtrate  from  Group  4  should  be  boiled  until  all  the 
in  the  solution  has  been  decomposed.  Add  to  the 
solution  enough  NH4OH  to  give  a  strong  alkaline  reaction.  If 
a  precipitate  should  form  with  the  NH4OH,  add  NH4C1  until  it 
dissolves.  The  group  is  now  precipitated  with  an  excess  of 
ammonium  carbonate  (NH4)2CO3,  boiled  for  a  moment,  and  fil- 
tered, care  being  taken  that  an  excess  of  NH4OH  is  always 
present.  The  carbonates  thus  precipitated  are  somewhat  soluble 


EXAMINATION  FOR   METALS  IN  SOLUTION  97 

in  ammonium  chlorid,  which  is  always  present  in  this  solu- 
tion from  the  preceding  groups.  The  boiling  tends  to  make 
the  precipitate  crystalline  and  less  soluble  in  the  NH4C1.  The 
filtrate  is  set  aside  for  further  examination,  the  precipitate  con- 
taining all  the  metals  of  this  group  which  may  be  present.  This 
precipitate  need  not  be  washed. 

Make  a  preliminary  test  to  see  if  barium  is  present  by  taking 
a  small  portion  of  the  precipitate,  dissolving  it  in  a  drop  or 
two  of  acetic  acid,  and  adding  potassium  chromate,  K2CrO4. 
A  yellow  precipitate  indicates  barium. 

If  barium  is  found  to  be  present,  dissolve  the  whole  precipi- 
tate in  acetic  acid,  add  K2CrO4  in  a  sufficient  excess  to  make  the 
whole  solution  decidedly  yellow,  filter,  and  wash.  All  the  barium 
is  precipitated,  while  the  strontium  and  calcium  remain  in  the 
filtrate.  Confirm  the  presence  of  barium  by  dissolving  some  of 
the  yellow  precipitate  in  HC1,  diluting  with  water,  and  adding 
some  dilute  H2SO4,  when  white  BaSO4  will  be  precipitated. 

The  pale  green  color  which  a  barium  compound  imparts  to  the 
non-luminous  flame  may  also  be  used  as  a  confirmatory  test. 

The  strontium  and  calcium  remaining  in  the  filtrate  are 
again  precipitated  with  (NH4)2CO3,  after  adding  an  excess  of 
NH4OH,  the  whole  boiled  for  a  moment,  filtered,  and  washed 
until  all  the  yellow  color  disappears  from  the  precipitate,  and 
the  filtrate  comes  through  colorless.  This  last  filtrate  may  be 
thrown  away. 

The  last  precipitate,  containing  the  carbonates  of  strontium 
and  calcium,  or  the  original  group  precipitate  if  barium  is 
absent,  is  then  dissolved  in  dilute  HC1. 

Make  a  preliminary  test  to  see  if  strontium  is  present,  by 
taking  a  small  portion  of  this  solution  and  adding  a  little  calcium 
sulfate,  CaSO4.  A  white  precipitate  shows  the  presence  of 
strontium.  This  preliminary  test  should  be  thrown  away. 

[The  coloration  of  the  non-luminous  flame  is  often  employed 
as  a  preliminary  test  to  show  the  presence  of  strontium.  Since 


98  QUALITATIVE  ANALYSIS 

strontium  and  calcium  both  give  colored  flames,  the  former  a 
carmine  red  and  the  latter  an  orange  or  brick  red,  unless  great 
care  is  used  there  is  danger  of  error.  A  comparison  of  the  two 
colored  flames  will  easily  show  the  difference.  It  is  therefore 
best,  especially  for  one  using  this  test  who  is  not  perfectly 
familiar  with  the  two  colors,  to  compare  the  color  of  the  flame 
of  the  unknown  substance  with  that  produced  by  calcium 
chlorid.] 

If  strontium  is  found  to  be  present,  add  a  dilute  solution 
of  ammonium  sulfate  (NH4)2SO4.  [A  solution  of  the  proper 
strength  is  made  by  dissolving  two  grams  of  (NH4)2SO4  in 
100  cc.  of  water.]  The  precipitation  takes  place  slowly  and  is 
complete  only  after  standing  some  time.  This  precipitates  the 
strontium  as  sulfate,  while  the  calcium  sulfate  remains  in  the 
dilute  solution,  it  being  quite  soluble  in  water. 

The  calcium  remaining  in  the  solution,  which  if  it  is  not  already 
so  should  be  made  alkaline  with  NH4OH,  is  then  precipitated 
by  ammonium  oxalate  as  white  calcium  oxalate,  CaC2O4. 

[If  too  large  an  excess  of  NH4C1  is  found  in  the  solution 
after  precipitating  the  third  and  fourth  groups,  it  will  interfere 
with  the  complete  precipitation  of  the  fifth  group,  the  metals  of 
which  will  then  be  precipitated  in  the  sixth  group  with  any 
magnesium  which  may  be  present.  To  avoid  this  the  filtrate 
from  Group  4  may  be  evaporated  to  dryness,  ignited,  and  the 
ammonium  compounds  volatilized  by  heat.  The  residue  is  dis- 
solved in  water,  adding  a  drop  or  two  of  HC1  if  necessary.  The 
solution  is  then  made  alkaline  with  NH4OH,  and  the  group 
precipitated  with  (NH4)2CO3.] 

Group  6 

To  this  group  belong  Magnesium,  Potassium,  Sodium,  and  the 
compound  radical  Ammonium. 

There  is  no  special  group  reagent  and  there  are  few  insoluble 
compounds  except  those  of  magnesium.  This  group  therefore 


EXAMINATION  FOR  METALS  IN  SOLUTION  99 

includes  all  metals  not  precipitated  by  the  preceding  group 
reagents. 

On  account  of  the  difficulty  in  making  a  complete  determina- 
tion of  the  alkali  metals,  potassium  and  sodium,  they  are  not 
always  sought  for  in  a  qualitative  analysis,  especially  since 
their  determination  in  a  quantitative  analysis  occurs  after  all 
the  other  metals,  including  magnesium,  have  been  removed, 
and  when  the  solution  can  contain  nothing  else. 

If  it  is  desirable  to  determine  their  presence  or  absence,  the 
filtrate  from  Group  5  must  be  subjected  to  a  preliminary  test 
at  this  point.  Their  presence  is  shown  by  the  coloration  which 
they  give  to  the  non-luminous  flame. 

The  flame  test  for  sodium  is  so  delicate  that  unless  the  char- 
acteristic yellow  flame  is  quite  brilliant  it  is  probably  due  to  an 
accidental  impurity. 

The  violet  flame,  due  to  potassium,  is  entirely  concealed  by 
the  yellow  flame  if  sodium  is  present.  It  can  then  be  seen 
through  a  blue  glass,  which  cuts  off  all  the  yellow  light.  [See 
Part  II,  IV.] 

If  sodium  is  absent,  the  filtrate  from  Group  5  is  made  strongly 
alkaline  with  NH4OH.  If  a  precipitate  should  occur  on  the 
addition  of  the  NH4OH,  add  NH4C1  until  the  precipitate  dis- 
solves. Add  acid  sodium  phosphate,  Na2HPO4,  in  excess,  and 
allow  the  whole  to  stand  for  some  hours  in  a  warm  place.  The 
magnesium  is  all  precipitated  as  white  crystalline  MgNH4PO4, 
which  is  soluble  in  acetic  acid. 

[It  sometimes  happens  that  some  metals  of  the  preceding 
groups,  particularly  aluminum  and  the  fifth-group  metals,  have 
not  been  completely  precipitated.  These  will  be  precipitated 
here  as  phosphates,  and  appear  at  once  as  a  white  flocculent 
precipitate.  If  this  precipitate  is  filtered  at  once,  enough  of 
the  magnesium  will  remain  in  the  filtrate  to  appear  after  a  time 
as  a  white  crystalline  precipitate,  adhering  more  or  less  closely 
to  the  sides  of  the  beaker.  This  can  always  be  seen  if  the 


100  QUALITATIVE  ANALYSIS 

inside  of  the  beaker  is  rubbed  with  a  glass  rod,  the  crystals 
forming  where  the  glass  is  rubbed.] 

If  sodium  is  present,  the  magnesium  must  be  precipitated  by 
some  other  reagent  to  avoid  introducing  sodium  into  the  solu- 
tion. This  may  be  done  by  .ammonium  phosphate,  or,  better, 
by  a  saturated  solution  of  barium  hydroxid,  Ba(OH)2.  The 
reagent  is  added  in  excess,  the  precipitate  filtered  off,  and 
the  excess  of  Ba(OH)2  removed  by  (NH4)2CO3.  The  barium 
carbonate  is  then  filtered  off,  the  filtrate  evaporated  to  dry- 
ness  in  a  porcelain  dish,  and  cautiously  ignited  to  expel  the 
ammonium  salts. 

The  residue,  containing  the  potassium  and  sodium  salts,  is 
dissolved  in  a  very  small  quantity  of  water,  filtered,  if  necessan-, 
and  divided  into  two  portions. 

To  one  portion  add  an  excess  of  acid  sodium  tartrate, 
NaH(C4H4O6),  and  allow  it  to  stand  for  some  time  with  occa- 
sional shaking.  A  white  crystalline  precipitate  of  KH(C4H4O6) 
indicates  potassium. 

[A  better  test,  although  not  often  employed  in  a  qualitative 
analysis  on  account  of  the  cost  of  the  reagent,  is  to  add  a  few 
drops  of  platinum  chlorid,  PtCl4,  and  evaporate  to  dry  ness  on 
a  water  bath.  The  residue  is  then  moistened  with  water,  and  a 
little  alcohol  added.  A  heavy,  yellow  crystalline  precipitate, 
insoluble  in  alcohol,  is  potassium  chloroplatinate,  K2PtCl6,  and 
indicates  potassium.] 

To  the  other  portion  add  acid  potassium  pyroantimonate, 
K2H2Sb2O7,  and  allow  it  to  stand  some  time.  A  white  precipi- 
tate indicates  sodium. 

The  presence  of  ammonium  is  shown  by  adding  concentrated 
NaOH  to  the  original  solution  and  warming.  Any  ammonium 
compound  present  is  decomposed,  and  ammonia,  NH3,  is  liber- 
ated, which  is  recognized  by  its  characteristic  odor. 


SYSTEMATIC  EXAMINATION  FOR  ACID  RADICALS 
IN  SOLUTION 

PRELIMINARY  EXAMINATION 

We  have  already  noted  (page  75)  that  there  is  no  simple 
method  for  the  determination  of  the  acid  radicals  by  successive 
elimination,  as  there  is  in  the  examination  for  the  metals. 
Fortunately,  in  most  cases  only  a  few  acids  need  be  looked  for 
in  mixed  compounds  in  solution,  since  the  presence  of  any 
metals,  except  those  of  the  alkali  group,  proves  conclusively 
the  absence  of  one  or  more  of  the  acid  radicals. 

As  an  illustration  of  this  fact  we  may  note  that  if  silver  is 
found  in  an  acid  solution,  it  would  be  impossible  for  any  hydro- 
chloric acid  to  be  present ;  or,  if  a  solution  contains  barium,  sul- 
furic  acid  could  not  be  present ;  or,  if  a  neutral  solution  contains 
lead,  only  nitric  and  acetic  need  be  looked  for,  since  all  other  lead 
salts  are  insoluble  in  water.  If  the  student  will  bear  in  mind 
the  above  examples,  and  many  similar  ones  which  will  readily 
suggest  themselves,  he  will  be  saved  much  useless  labor. 

The  first  step  in  analyzing  a  solution  is,  therefore,  to  deter- 
mine what  metals  are  present.  Having  done  this  we  know 
that  all  those  acids  which  form  insoluble  compounds  with  any 
of  the  metals  which  are  present  must  be  absent,  and  we  may 
then  proceed  to  determine  the  acids  that  are  present.  [The 
word  "acid,"  as  used  in  this  connection,  does  not  necessarily 
mean  free  acid,  but  is  used  to  denote  the  acid  radical  which  is 
combined  with  a  metal  to  form  the  salt.] 

A  number  of  salts  which  are  insoluble  in  water  are  kept  in 
solution  by  the  presence  of  some  free  acid.  When  the  latter  is 
neutralized  in  the  examination  for  the  metals,  not  only  the 

101 


102  QUALITATIVE   ANALYSIS 


but  often:  4he:  acid  will  be  precipitated  ;  so  that  a  number 
of  the  acids,  if  present,  will  be  discovered  during  the  examina- 
tion /on  the  njejbalsr  •  -  -Thus,  if  phosphoric,  oxalic,  silicic,  boric, 
or  hydrofluoric  acid  is  present,  together  with  certain  members 
of  Groups  3,  4,  5,  or  6  (Mg),  it  will  remain  in  solution  as 
long  as  free  acids  are  present,  but  will  be  precipitated  with 
Group  3,  and  can  be  detected  as  given  on  page  90. 

PREPARATION   OF   THE   SOLUTION 

First  Method.  After  having  determined  what  metals  are  pres- 
ent, a  solution  must  be  carefully  prepared  before  examining  for 
the  acids.  For  this  purpose  take  a  portion  of  the  original  solu- 
tion (50  cc.)  and  heat  nearly  to  boiling.  Add  to  the  hot  solution 
a  very  slight  excess  of  Na2CO3,  boil  for  a  moment,  and  filter. 
Preserve  some  of  the  precipitate  for  a  future  test. 

The  filtrate,  which  contains  all  the  acid  radicals,  and  which 
should  be  alkaline  from  the  excess  of  Na2CO3,  must  now  be 
neutralized,  and  the  excess  of  the  sodium  carbonate  removed. 
In  order  to  do  this,  while  the  liquid  is  kept  boiling  add  HNO3 
carefully  and  slowly,  at  the  last  drop  by  drop,  as  long  as  the 
carbon  dioxid  continues  to  escape. 

After  the  carbonate  is  thus  decomposed,  boil  for  a  moment, 
and  then  test  with  litmus  paper.  The  liquid  should  now  be 
slightly  acid.  Next  add  to  the  liquid  dilute  NH4OH  drop  by 
drop,  stirring  all  the  while,  until  it  is  exactly  neutral.  If  very 
much  of  an  excess  of  acid  has  been  added,  it  is  better  to  very 
nearly  neutralize  this  with  NaOH,  completing  the  operation 
with  the  NH4OH.  The  reason  for  this  is  that  the  presence  of 
ammonium  salts  in  any  but  minute  quantities  interferes  with 
the  precipitation  of  some  of  the  acid  radicals.  Great  care 
should  therefore  be  taken  in  the  preparation  of  this  solution, 
since  much  of  the  success  in  the  determination  of  the  acid  radi- 
cals depends  upon  the  care  with  which  this  solution  has  been 
prepared. 


EXAMINATION  FOR  ACID  RADICALS  IN  SOLUTION      103 

Second  Method.  If  only  members  of  the  first  and  second 
groups  of  the  metals  were  found  in  the  original  solution, 
they  may  be  removed  by  means  of  H2S  instead  of  Na2CO3. 

To  prepare  the  solution  by  this  method,  saturate  the  original 
solution  with  H2S,  heat  the  whole  to  boiling,  and  filter.  Boil 
the  filtrate  until  all  the  H2S  has  been  expelled.  The  solution 
should  now  be  slightly  acid.  Carefully  neutralize  the  acid  as 
in  the  first  method,  and  the  solution  is  ready  for  use. 

Arsenic,  when  present,  usually  belongs  with  the  acid  radical. 
In  the  examination  for  the  metals,  arsenic  acid,  when  present,  is 
reduced  to  arsenious  acid  by  H2S.  [See  Part  I,  Arsenic.] 

In  the  preparation  of  the  solution  for  acids  by  the  first 
method  the  arsenic  is  not  precipitated  by  the  Na2CO3,  and  so, 
if  found  in  the  examination  for  metals,  may  now  be  looked  for 
among  the  acids. 

If  the  second  method  for  preparing  the  solution  for  acids  is 
employed,  the  arsenic  will  be  precipitated  with  the  metals  and 
so  will  not  be  found  among  the  acids. 

CLASSIFICATION  OF  THE  ACID  RADICALS 

The  classification  of  the  acid  radicals  is  based  upon  the  solu- 
bility or  insolubility  of  certain  of  their  salts,  which  are  pre- 
cipitated by  certain  reagents.  They  are  thus  divided  into 
certain  arbitrary  groups,  as  in  the  case  of  the  metals,  but  the 
methods  of  treatment  for  these  acid  groups  are  quite  different. 
These  groups  can  be  subdivided  only  to  a  limited  extent,  so 
that  in  most  cases  each  acid  has  to  be  detected  by  a  special 
test.  The  group  reagents  are  therefore  employed  to  ascertain 
the  presence  or  absence  ot  any  of  the  members  of  a  group  rather 
than  to  separate  them  from  one  another. 

In  the  classification  of  the  acid  radicals  we  recognize  three 
main  groups,  the  grouping  being  based  upon  the  insolubility  of 
their  barium  and  silver  salts.  Only  the  comparatively  common 
acids  will  be  treated  here.  For  the  detection  of  the  rare  acids 


104  QUALITATIVE   ANALYSIS 

the  student  should  consult  some  large  work  like   Fresenius' 
Manual  of  Qualitative  Chemical  Analysis. 

Group  1  contains  those  acid  radicals  whose  barium  salts  are 
precipitated  from  a  neutral  solution  by  means  of  barium  chloricl. 

Group  2  contains  those  acid  radicals  whose  silver  salts  are 
precipitated  from  a  solution  acidified  with  nitric  acid  by  means 
of  silver  nitrate. 

Group  3  contains  the  remaining  acid  radicals ;  that  is,  those 
whose  barium  salts  are  soluble  in  water,  and  whose  silver  salts 
are  soluble  in  water  or  in  nitric  acid,  or  in  both. 

Acids  :  Group  I 

The  neutral  solution  prepared  for  the  examination  for  acids 
as  described  above  is  treated  with  barium  chlorid,  filtered,  and 
the  filtrate  preserved  for  some  tests  in  the  next  group.  The 
barium  salts  of  the  following  acids  may  be  found  in  the 
precipitate,  viz. : 

Sulfuric  acid,  H2S04, 
Sulfurous  acid,  H2S03, 
Phosphoric  acid,  H3P04, 
Ajsenic  acid,  H3As04, 
Arsenious  acid,  H3As03, 
Boric  acid,  H3B03, 
Hydrofluoric  acid,  HF, 
Carbonic  acid,  H2C03, 
Silicic  acid,  H4Si04  or  H2Si03, 
Chromic  acid,  H2Cr04, 
Oxalic  acid,  H2C204, 
Tartaric  acid,  H2(C4H406),  and 
Citric  acid,  H3(C6H507). 

Take  a  portion  of  the  precipitate  and  add  to  it  some  HC1. 
If  it  remains  undissolved,  sulfuric  acid  is  present.  If  sulfur 
dioxid  is  given  off  (recognized  by  its  odor),  sulfurous  acid  is 


EXAMINATION  FOR  ACID  RADICALS  IN  SOLUTION      105 

present.  The  sulfites  should  completely  dissolve  in  HC1,  but 
as  they  gradually  oxidize  to  sulfates,  even  by  standing,  there 
is  usually  a  small  portion  which  will  not  dissolve.  [See  Part  I, 
Sulfurous  Acid  3.] 

If  carbonates  are  present,  they  dissolve  in  HC1  with  effer- 
vescence, which  is  caused  by  the  liberation  of  carbon  dioxid. 
The  latter  can  be  recognized  by  the  ordinary  test  with  lime- 
water.  [See  Part  I,  Carbonic  Acid  4.] 

If  members  of  Groups  3,  4,  5,  or  6  (Mg)  of  the  metals  were 
found  in  the  examination  for  metals,  phosphoric,  oxalic,  hydro- 
fluoric, boric,  and  silicic  acids,  if  present,  were  probably  found 
in  the  precipitate  with  Group  3,  and  recognized  by  the  tests 
given  there.  [See  page  90.]  In  any  case  if  it  is  desirable  to 
test  for  them  here,  the  tests  there  given  are  the  ones  to  be 
employed. 

If  arsenic  or  arsenious  acid  is  present,  the  arsenic  acid  will 
be  reduced  to  arsenious  acid  by  H2S,  and  both  will  be  precipi- 
tated in  Group  2  of  the  metals. 

If  chromic  acid  is  present,  it  will  have  been  reduced  by  H2S 
and  precipitated  by  the  ammonium  hydroxid  in  Group  3  of  the 
metals.  In  this  acid  separation  the  presence  of  chromic  acid 
imparts  a  yellow  color  to  the  solution,  and  its  presence  may  be 
confirmed  by  making  the  solution  strongly  acid  with  acetic 
acid  and  adding  lead  acetate.  A  yellow  precipitate  soluble 
in  NaOH  is  lead  chromate.  Chromium  may  also  have  been 
present  in  the  original  solution  as  a  chromium  salt  of  some  acid. 
To  determine  whether  or  not  this  is  so,  take  a  small  portion  of 
the  precipitate,  obtained  by  boiling  the  original  solution  with 
Na2CO3,  and  dissolve  it  in  HC1.  Precipitate  this  last  solution 
with  NH4OH  or  NaOH.  Filter,  wash  the  precipitate,  and  test 
it  for  chromium  as  given  on  page  89. 

If  a  non-volatile  organic  acid,  such  as  tartaric  or  citric  acid, 
is  present,  some  of  the  metals  will  not  be  completely  precipi- 
tated in  the  preparation  of  the  solution  for  acids.  [See  Part  I, 


106  QUALITATIVE  ANALYSIS 

Aluminum  2,  3,  4,  and  Ferric  1.]  The  presence  of  these 
metals  may  interfere  with  the  test  for  these  a.cids,  and  so,  if 
present,  they  must  be  removed.  In  order  to  remove  them,  take 
some  of  the  solution  prepared  for  the  examination  for  acids, 
make  it  alkaline  with  NH4OH,  saturate  it  with  H2S,  heat  it  to 
boiling,  and  filter.  Make  the  filtrate  slightly  acid  with  HC1 
and  boil  until  all  the  H2S  has  been  expelled.  Filter  again  if 
necessary.  This  will  remove  any  metals  remaining  in  the 
solution  except  traces  of  aluminum  and  chromium,  and  they 
will  do  no  harm. 

After  the  solution  is  cold,  add  NH4OH  to  distinct  alkaline 
reaction,  then  NH4C1,  and  finally  CaCl2.  Shake  or  stir  vigor- 
ously and  allow  to  stand  fifteen  minutes.  If  no  precipitate 
appears,  tartaric  acid  is  absent.  If  a  precipitate  appears,  it  may 
or  may  not  be  caused  by  tartaric  acid,  but  the  solution  should 
be  allowed  to  stand  two  hours  for  complete  precipitation.  After 
this,  filter  and  save  the  filtrate  to  test  for  citric  acid.  The  pre- 
cipitate is  dissolved  in  cold  NaOH,  allowed  to  stand  a  few 
minutes  with  occasional  stirring,  filtered,  and  the  filtrate  boiled. 
If  a  precipitate  separates  by  boiling,  it  indicates  tartaric  acid, 
which  may  now  be  filtered  off  and  tested  as  under  Part  I, 
Tartaric  Acid  5. 

Citric  acid  is  not  very  often  found,  but,  if  present,  it  may  be 
detected  in  the  filtrate  after  precipitating  the  tartrates  with 
CaCl2.  To  do  this,  add  to  this  filtrate  three  volumes  of  alco- 
hol, and  after  allowing  it  to  stand  for  a  moment,  filter  and  wash 
the  precipitate  with  alcohol.  Dissolve  the  precipitate  on  the 
filter  with  a  little  dilute  HC1,  add  NH4OH  to  this  solution 
until  alkaline,  and  boil.  If  no  precipitate  appears,  add  a  little 
more  CaCl2  and  NH4OH  and  boil  again.  A  precipitate  is 
calcium  citrate  and  indicates  citric  acid. 


EXAMINATION  FOR  ACID  RADICALS  IN  SOLUTION      107 


Acids  :  Group  2 

To  a  second  portion  of  the  prepared  solution  add  dilute  HNO3 
until  distinctly  acid,  and  boil.  This  decomposes  sulfids,  sulfites, 
thiosulfates,  and  nitrites.  These  give  off  H2S,  SO2,  and  N2O3, 
respectively,  which  serves  to  recognize  them.  If  silver  nitrate 
is  now  added,  the  silver  salts  of  the  following  acids  may  be 
precipitated. 

Hydrochloric  acid,  HC1, 
Hydrobromic  acid,  HBr, 
Hydriodic  acid,  HI, 
Hydrocyanic  acid,  HCN, 
Hydroferrocyanic  acid,  H4Fe(CN)6, 
Hydroferricyanic  acid,  H8Fe(CN)6, 
Sulfocyanic  or  Thiocyanic  acid,  HSCN. 

In  addition  to  these,  phosphoric  and  chromic  acids,  if  present, 
will  be  precipitated  again  in  this  group,  unless  too  much  HNO3 
has  been  added. 

If  hydrogen  sulfid  was  present,  unless  it  was  completely 
decomposed  when  boiled  with  HNO3,  the  precipitate  will  be 
black,  in  which  case  add  a  little  more  dilute  HNO3,  boil  until 
all  the  H2S  is  expelled,  and  filter. 

Since  the  presence  of  any  of  the  cyanogen  acids  interferes 
with  the  tests  for  the  halogen  acids,  they  must  first  be  removed. 
The  presence  or  absence  of  the  cyanogen  acids  is  determined  in 
the  following  way. 

Take  a  little  of  the  prepared  solution,  add  a  few  drops  of 
HC1  and  some  ferric  chlorid.  A  red  solution  indicates  a  sulfo- 
cyanate,  a  blue  precipitate,  a  ferrocyanid.  If  both  are  present, 
the  blue  color  covers  up  the  red,  in  which  case  add  to  the 
solution  about  2  cc.  of  ether  and  shake  thoroughly.  The 
sulfocyanate,  if  present,  dissolves  in  the  ether  and  gives  it  a 
red  color. 


108  QUALITATIVE  ANALYSIS 

Take  a  little  of  the  filtrate  from  the  first  acid  group,  add  a 
few  drops  of  HNO3,  and  then  some  AgNO3.  A  red  brown  pre- 
cipitate indicates  a  ferricyanid.  [This  may  also  be  shown  by 
adding  to  the  solution  before  preparing  it  for  the  acid  tests 
some  freshly  prepared  solution  of  ferrous  sulfate,  when  Turn- 
bull's  blue  is  formed.] 

If  neither  hydroferrocyanic  nor  hydroferricyanic  acid  is 
present,  take  a  little  of  the  prepared  solution,  acidify  with 
HC1,  and,  while  cold,  add  a  few  drops  of  both  FeSO4  and 
FeClg.  If  the  solution  is  acid,  there  should  be  no  precipitate. 
Now  add  NaOH  until  the  mixture  is  strongly  alkaline,  warm  a 
little,  and  add  HC1  until  acid.  A  blue  precipitate  indicates 
hydrocyanic  acid. 

If  either  hydroferrocyanic  or  hydroferricyanic  acid  is  found 
to  be  present,  place  a  small  portion  of  the  precipitate,  formed  by 
AgNOg  in  precipitating  acid  Group  2,  in  a  small  dish,  add  to 
it  a  little  dilute  H2SO4,  and  warm.  Hydrocyanic  acid  will  be 
liberated  and  may  be  recognized  by  its  odor.  [Caution !  This 
latter  operation  should  be  carried  on  under  a  hood  with  a  good 
draught,  and  the  liberated  gas  smelled  cautiously,  as  it  is  very 
poisonous.] 

If  any  of  the  cyanogen  acids  are  found  in  the  examination 
as  given  above,  they  must  now  be  removed.  This  can  be  done 
in  the  following  manner. 

Take  about  one  half  of  the  group  precipitate  formed  by 
AgNOg,  place  it  in  a  porcelain  crucible,  and  ignite  for  about 
five  minutes  over  a  Bunsen  lamp.  The  silver  salts  of  the 
cyanogen  acids  will  be  decomposed,  while  the  silver  salts  of 
the  halogen  acids  will  remain  unchanged.  Add  to  the  ignited 
residue  about  four  times  its  weight  of  NaKCO3  and  thoroughly 
fuse  the  whole.  The  fused  mass  is  then  digested  for  some  time 
with  warm  water  and  filtered.  The  chlorids,  bromids,  and  iodids 
will  now  be  found  in  the  filtrate  as  alkaline  salts,  free  from 
cyanogen  compounds. 


EXAMINATION  FOR  ACID  RADICALS  IN  SOLUTION     109 

The  test  for  hydrochloric  acid  in  the  presence  of  hydrobromic 
and  hydriodic  acids  requires  some  care,  but  it  may  be  done  as 
follows,  with  the  solution  just  prepared. 

Take  about  10  cc.  of  the  solution  in  a  large  test-tube  and 
make  it  distinctly  acid  with  H2SO4.  Boil  to  decompose  the 
excess  of  Na2CO3.  Add  about  3  cc.  of  a  concentrated  solution 
of  ferric  sulfate  and  boil  again.  [Any  of  the  double  sulfates, 
known  as  the  iron  alums,  may  be  used.  Ferric  chlorid  may  also 
be  used  if  perfectly  free  from  nitrates.]  Test  the  escaping  steam 
for  iodin  with  starch  paper,  and,  if  found,  boil  until  the  iodin  is 
all  expelled,  adding  more  of  the  iron  solution  if  necessary. 

After  the  iodin  is  all  expelled,  add  a  solution  of  KMnO4  to 
distinct  coloration  and  boil  again.  Test  the  escaping  steam  for 
bromin  with  potassium  iodid  starch  paper.  [Do  not  allow  the 
liquid  to  touch  the  paper.]  If  bromin  is  found,  boil  until  it  is 
expelled,  adding  enough  of  the  KMnO4  so  that  the  solution  will 
show  the  purple  color  after  boiling. 

The  excess  of  the  KMnO4  is  removed  by  adding  a  few  drops 
of  alcohol,  boiling  a  moment,  and  filtering.  The  presence  of 
chlorin  can  now  be  shown  by  adding  a  little  HNO3  and  some 
AgNO3.  A  white  precipitate  is  silver  chlorid.  [Instead  of 
KMnO4  a  solution  of  K2Cr2O7  may  be  used.] 

If  none  of  the  cyanogen  acids  were  found  to  be  present,  the 
halogen  acids  may  be  tested  for  in  the  following  way. 

Take  some  of  the  group  precipitate  formed  by  AgNO3  and 
place  it  in  a  porcelain  dish  with  some  dilute  H2SO4  and  a  piece 
of  metallic  zinc.  The  nascent  hydrogen-reduces  the  silver  salts 
to  metallic  silver,  which  may  be  filtered  off.  The  excess  of 
zinc  is  removed  with  Na2CO3,  as  in  the  preparation  of  the  solu- 
tion for  acids.  The  alkaline  filtrate  is  then  made  acid  with 
H2SO4,  which  decomposes  the  excess  of  Na2CO3.  The  solution 
may  now  be  examined  for  the  halogen  acids  as  above. 

If  nitrous  acid  was  present  in  the  original  solution,  it  will 
have  been  decomposed  by  the  nitric  acid  in  the  preparation  of 


110  QUALITATIVE   ANALYSIS 

the  solution  for  this  group.  Its  presence  may  be  confirmed  by 
adding  acetic  or  dilute  sulfuric  acid  to  the  original  solution, 
when  red-brown  fumes  of  N2O3  or  NO2  will  be  liberated. 

Acids:  Group  3 
The  only  acid  radicals  found  in  this  group  are 

Nitric  acid,  HN08, 
Chloric  acid,  HC103,  and 
Acetic  acid,  H(C2H302). 

These  cannot  be  precipitated  under  ordinary  conditions  and 
so  are  recognized  by  special  tests. 

Nitric  acid  may  be  recognized  in  the  original  solution  by 
either  of  the  tests  already  given.  [See  Part  I,  Nitric  Acid  1 
and  2.] 

Nitrous  acid  shows  the  same  result  as  in  the  first  test  for  nitric 
acid  referred  to  above.  If,  however,  a  portion  of  the  solution 
is  acidified  with  dilute  H2SO4  and  boiled,  all  nitrites  will  be 
decomposed  and  N2O3  or  NO2  liberated.  Any  other  dilute  acid 
will  produce  the  same  result  as  H2SO4. 

Chloric  acid  is  not  common  in  solutions..  When  present  it 
may  be  recognized  by  evaporating  a  small  quantity  of  the  pre- 
pared solution  to  dryness  in  a  porcelain  dish  on  the  water 
bath,  and  adding  to  the  residue  a  small  quantity  of  concen- 
trated H2SO4.  The  heavy,  yellow,  explosive  gas,  chlorin  per- 
oxid,  C1O2,  is  liberated,  which  may  be  further  recognized  by  its 
peculiar  "  chlorous  "  odor. 

Acetic  acid  may  be  recognized  by  its  special  test.  [See  Part 
II,  VIII,  (/).]  If  a  chlorate  is  present,  it  should  be  entirely 
decomposed  by  sulfuric  acid,  as  given  above,  before  making 
tfris  test. 


PART  IV 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS 

Hitherto  we  have  had  for  our  study  the  analysis  of  substances 
already  in  solution,  or,  as  in  Part  II,  the  study  of  simple  com- 
pounds in  the  solid  state.  If  the  analytical  chemist  could 
always  receive  his  complex  substances  in  solution,  his  work 
would  be  comparatively  simple,  and  the  methods  given  in 
Part  III  would  enable  him  to  analyze  all  ordinary  inorganic  sub- 
stances. As  a  matter  of  fact,  most  of  the  material  sent  to  the 
chemist  for  analysis  is  more  or  less  complex  and  usually  in  the 
solid  state,  and  it  is  necessary  for  him  to  know  how  to  proceed 
under  these  conditions. 

It  often  happens  that  the  physical  properties  of  a  substance 
are  such  that  the  chemist  is  able  to  infer  many  things  as  to  its 
nature  and  composition.  His  knowledge  of  the  more  important 
reagents  and  their  action  upon  the  different  elements  and 
compounds  should  enable  him  to  determine  by  a  few  simple 
tests  whether  or  not  his  inference  is  correct. 

It  sometimes  happens,  also,  that  in  a  complex  body  the  only 
knowledge  desired  is  regarding  the  presence  or  absence  of  a 
single  ingredient.  The  method  of  procedure  in  such  a  case  is 
not  very  complicated  and  the  result  is  easily  acquired  by 
methods  already  given. 

But  if  we  wish  to  know  all  the  ingredients  to  be  found  in  a 
complex  chemical  compound  or  mixture  ;  if  we  wish  to  know 
that  it  contains  certain  ingredients  and  no  others ;  in  short, 
if  we  wish  to  make  a  complete  qualitative  analysis  of  the 

111 


112  QUALITATIVE  ANALYSIS 

substance,  we  must  employ  a  systematic  course  of  analysis.  We 
must  of  course  know  the  common  reagents  and  their  action  on 
all  kinds  of  matter,  but  haphazard  testing  will  not  accomplish 
our  purpose.  We  must  know  what  solvents  and  reagents  to 
use  and  in  what  order  they  must  be  applied,  and  follow  this 
order,  or  our  analysis  will  be  little  better  than  guess  work. 

A  preliminary  examination  of  the  substance  should  always 
be  made,  since  that  will  often  reveal  the  absence  of  some  ingre- 
dient which  will  enable  the  chemist  to  shorten  his  work,  or  the 
presence  of  some  ingredient  which  will  enable  him  to  take  cer- 
tain precautions,  and  so  avoid  some  difficulty.  In  any  case,  the 
preliminary  examination  will  always  indicate  the  nature  of  the 
substance  to  be  analyzed. 

PRELIMINARY  EXAMINATION 

Since  the  object  of  a  preliminary  examination  is  to  obtain 
information  regarding  the  nature  of  the  substance  to  be  ana- 
lyzed, considerable  care  should  be  exercised  in  considering  any 
physical  or  chemical  properties  which  may  be  observed,  as  this 
may  result  in  the  saving  of  much  time  and  labor. 

The  substance  should  first  be  examined  with  regard  to  its 
physical  properties.  Note  its  color  and  structure,  whether  it 
is  homogeneous  or  heterogeneous,  crystalline  or  amorphous ; 
or,  if  it  is  a  compact  mass,  whether  it  is  hard  or  soft;  or,  if 
metallic,  whether  it  is  malleable  or  brittle,  and  whether  it  has 
a  high  or  low  specific  gravity.  These  are  the  principal  phys- 
ical properties  to  note,  and  the  student's  knowledge  of  these 
properties  should  give  him  considerable  information  about  the 
substance. 

Most  of  the  preliminary  tests  employed  to  show  the  chemical 
properties  are  included  in  those  reactions  which  are  given 
under  Blowpipe  Analysis  (Part  II),  and  consist  of  any  or  all  of 
the  different  methods  of  treatment  which  are  given  in  that 
part  of  the  work.  It  will  of  course  sometimes  happen  that  the 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS      113 

tests  given  there  will  have  little  or  no  value,  since  in  a  complex 
substance  two  elements  may  be  present  which  will  give  con- 
flicting results,  and  so  will  modify  those  results  which  we 
should  otherwise  obtain.  Many  of  the  reactions  will,  however, 
give  important  indications.  The  tests  given  in  Part  II  should 
be  followed  in  about  the  same  order  in  which  they  are  given 
there,  the  detail  of  which  need  not  be  repeated  here. 

The  methods  here  given  provide  for  the  recognition  of  only 
the  common  inorganic  substances  and  the  salts  of  a  few  of  the 
simplest  and  most  common  organic  acids.  If  other  organic 
material  is  found  to  be  present,  it  may  be  removed  by  igniting 
the  substance  in  a  crucible  until  all  the  organic  matter  has 
been  oxidized  and  driven  off. 

All  solid  substances  may  be  divided  into  two  general  classes, 
viz.  : 

I.    Metals  and  alloys. 

II.    Substances  which  are  neither  metals  nor  alloys, 

A  metal  or  an  alloy  can  usually  be  recognized  as  such  by  its 
metallic  luster,  hardness,  malleability,  and  high  specific  grav- 
ity. Only  a  few  of  the  metals  or  alloys  are  brittle,  and  only 
a  very  few  have  a  specific  gravity  of  less  than  four. 

Substances  which  are  neither  metals  nor  alloys  usually  lack 
the  metallic  luster,  are  more  or  less  brittle,  though  sometimes 
very  hard,  and  have  a  specific  gravity  less  than  four.  The  most 
of  them  can  be  reduced  to  a  powder  with  comparative  ease. 

I.    THE  SUBSTANCE  IS  A  METAL  OR  AN  ALLOY 

This  class  may  be  further  divided  into  three  divisions,  accord- 
ing to  the  action  of  the  different  substances  when  treated  with 
nitric  acid.  These  divisions  are  : 

A.  Metals  insoluble  and  unchanged  in  nitric  acid.  This  division 
includes  gold,  platinum  and  most  of  the  rare  metals  of  the 
platinum  group,  and  their  alloys. 


114  QUALITATIVE  ANALYSIS 

B.  Metals  which  form  insoluble  oxids  by  the  action  of  nitric  acid. 
This  division  includes  antimony  and  tin  and  their  alloys. 

C.  Metals  and  alloys  soluble  in  nitric  acid.     This  division  includes 
all  the  other  common  metals  and  their  alloys. 

Since  chemically  pure  metals  are  not  often  found,  and,  if 
found,  are  rarely  the  object  of  a  qualitative  analysis,  only  the 
general  process  of  treating  alloys  will  be  given  here,  since  that 
will  be  sufficient  for  all  cases. 

No  metals  used  commercially  are  ever  chemically  pure,  but 
all  contain  greater  or  less  quantities  of  other  metals  which 
have  not  been  completely  removed  in  the  process  of  reducing 
the  metals  from  their  ores.  These  metals  can  hardly  be  classed 
as  alloys,  and  yet,  for  analytical  purposes,  they  are  such.  If 
an  analysis  of  the  impurity  in  a  commercial  metal  is  desired,  it 
is  only  necessary  to  employ  a  larger  amount  and  proceed  as  in 
the  analysis  of  an  alloy. 

Heat  a  portion  (1  or  2  grams)  of  the  alloy  with  nitric  acid 
of  about  1.2  specific  gravity  (equal  parts  of  concentrated  acid 
and  water).  After  the  action,  if  any,  has  ceased,  if  a  white 
residue  is  left,  add  an  equal  volume  of  water  and  heat  to  boil- 
ing. [Certain  nitrates  are  not  very  soluble  in  nitric  acid,  and 
so  crystallize  out  if  the  solution  is  too  concentrated.]  Filter 
and  wash  the  residue  with  a  little  hot  water.  Test  the 
filtrate  by  evaporating  a  small  portion  on  platinum  foil  to  see 
if  anything  has  dissolved.  If  a  residue  remains  on  the  foil,  it 
belongs  to  division  C,  and  the  filtrate  must  be  preserved  and 
tested  as  given  under  C. 

A.  The  residue  is  dark  colored,  or  has  a  metallic  appearance. 
Dissolve  it  in  aqua  regia,  evaporate  to  dryness,  add  a  little 
concentrated  HC1,  and  evaporate  a  second  time  ;  after  which, 
add  some  water,  and  evaporate  a  third  time  to  remove  the  excess 
of  acid.  Dissolve  the  residue  in  water,  add  oxalic  acid,  and 
warm  for  an  hour,  when  all  the  gold  will  be  precipitated.  Filter 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS      115 

off  the  gold  and  evaporate  the  filtrate  to  dryness.  Ignite  the 
dried  filtrate  in  a  porcelain  crucible,  under  a  hood,  to  remove 
the  excess  of  oxalic  acid.  The  platinum  will  remain  in  the 
crucible  in  the  metallic  state,  and  may  be  dissolved  again  in 
aqua  regia  and  subjected  to  further  tests. 

B.  The  residue  is  white  and  pulverulent.     This  generally  consists 
of  metastannic  acid,  H10Sn5O15,  or  metantimonic  acid,  HSbO3.    If 
the  residue  is  dark  colored,  it  may  consist  of  both  divisions  A 
and  B.     Fuse  this  residue  in  a  porcelain  crucible,  with  a  mixture 
of  equal  parts  of  dry  Na2CO3  and  sulfur,  for  five  or  ten  minutes. 
After  the  crucible  and  its  contents  have  become  cool,  dissolve 
out  the  fused  mass  with  hot  water,  and  filter  if  necessary.     A 
precipitate  will  consist  of  division  A,  and  is  treated  as  given 
above.     The  solution,  or  filtrate,  should  now  be  acidified  with 
dilute  HC1,  which  precipitates  the  antimony  and  tin  as  sulfids, 
after  which  they  may  be  separated,  as  given  in  Part  III,  Group 
2,  Subdivision  B.     [See  page  84.] 

C.  The  substance  is  completely  dissolved  by  HN03.     The  filtrates 
from  A  and  B,  or  the  entire  solution  if  the  alloy  dissolved  com- 
pletely in  HNO3,  should  now  be  evaporated  to  dryness,  and 
then  dissolved  in  water,  a  few  drops  of  HNO3  being  added,  if 
necessary,  to  make  a  clear  solution. 

The  metals  are  now  in  solution  as  nitrates  and  may  be  sepa- 
rated by  the  methods  given  under  Part  III,  Mixed  Compounds. 
[See  page  76.] 

H.    THE  SUBSTANCE  IS  NEITHER  A  METAL  NOR  AN  ALLOY 

In  this  case  if  the  substance  is  a  mixture,  the  different  ingre- 
dients are  first  separated  from  each  other,  as  far  as  possible, 
according  to  their  solubility  or  insolubility  in  water  and  acids. 
If  a  residue  remains  which  is  insoluble  in  water  and  acids,  it 
is  subjected  to  a  special  course  of  treatment  as  given  under  E. 


116  QUALITATIVE  ANALYSIS 

If  the  substance  for  analysis  is  a  complex  compound,  it  may 
be  treated  in  the  same  systematic  way,  but  will  usually  be  found 
to  be  insoluble  in  water  and  acids  and  will  have  to  be  brought 
into  solution  by  fusion  with  the  alkaline  carbonates,  as  given 
below  under  E,  (c).  If  the  substance  is  shown  by  the  prelimi- 
nary examination  to  be  a  silicate,  and  nothing  else,  we  may  pass 
at  once  to  division  F,  which  treats  of  silicates. 

If  the  substance  is  not  already  in  that  condition,  it  should 
be  well  pulverized  before  beginning  the  analysis.  From  one  to 
three  grams  should  be  taken  for  analysis. 

A.  The  substance  is  partially  or  entirely  soluble  in  water.     Place 
the  substance  in  a  proper-sized  beaker  or  flask,  add  to  it  dis- 
tilled water,  and  heat  to  boiling.     Digest  the  substance  at  or 
near  the  boiling  point  for  a  few  minutes,  stirring  or  shaking 
frequently,  and  filter  if  it  has  not  all  dissolved.     If  the  sub- 
stance completely  dissolves,  the  solution  may  now  be  examined 
for  metals  as  given  in  Part  III,  page  76,  and  for  acid  radicals 
as  given  on  page  101. 

If  a  residue  remains  undissolved,  filter  as  above  directed, 
and  wash  it  thoroughly  with  hot  water.  Evaporate  a  small 
portion  of  the  filtrate  to  dryness  on  a  platinum  foil,  to  see  if 
anything  has  dissolved.  If  a  residue  remains  on  the  foil,  the 
whole  filtrate  must  be  examined,  as  directed  above,  when  all 
dissolves. 

B.  The  substance  is  insoluble  in  water.     A  small  portion  of  the 
undissolved  residue  from  A  is  subjected  to  a  preliminary  exami- 
nation with  warm  dilute  HC1.    This  may  be  done  in  a  test-tube. 
Note  carefully  if  any  gases  are  given  off,  as  these  indicate  what 
kind  of  compounds  are  being  dissolved.     A  carbonate  gives  off 
CO2  ;  a  sulfid  gives  off  H2S  ;  a  peroxid  or  some  highly  oxidized 
salt  (such  as  chromate,  etc.)  gives  off  Cl ;  a  sulfite  or  thiosulfate 
gives  off   SO2  ;    a  cyanid  gives   off  HCN   (Poison!).     These 
gases  may  be  recognized  by  the  tests  already  given.    Silicic  acid 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS       117 

sometimes  separates  out  as  a  white  gelatinous  mass,  indicating 
a  silicate  which  is  decomposed  by  acids.  [See  F.] 

If  no  action  takes  place,  test  a  few  drops  of  the  acid  liquid 
on  a  platinum  foil  to  see  if  anything  has  dissolved.  If  nothing 
has  dissolved,  throw  away  the  small  portion  taken  for  the  pre- 
liminary test  and  then  pass  on  to  C.  If  something  has  dis- 
solved, treat  the  whole  residue  from  A  with  HC1,  as  above. 
After  all  action  has  ceased,  boil  for  a  moment  and,  if  any  un- 
dissolved  residue  remains,  filter  and  wash  with  hot  water.  [The 
wash  water,  after  the  first  washing,  should  be  thrown  away.] 

The  residue  is  preserved  for  C  and  the  filtrate  evaporated  to 
dryness  in  a  porcelain  dish.  The  dried  mass  is  then  dissolved 
in  water  and  the  metals  separated,  as  in  Part  III. 

If  cyanids,  especially  the  compound  cyanids,  are  present, 
they  are  not  all  completely  dissolved  or  decomposed  by  the  HC1. 
They  will  therefore  be  found  in  some  of  the  successive  portions. 
A  more  complete  method  of  separation  under  these  conditions 
will  be  found  under  G. 

C.  The  substance  is  insoluble  in  1^0  and  in  HC1.     The  residue 
from  B  is  next  treated  with  warm  dilute  HNO3  (equal  parts 
concentrated  HNO3  and  water).     Observe  carefully  if  any  action 
takes  place.    If  nitric  oxid  is  given  off,  shown  by  the  red-brown 
fumes,  oxidation  is  taking  place.    Certain  sulfids  are  decomposed 
by  HNO3,  giving  off  some  H2S,  and  liberating  sulfur,  which 
appears  as  a  dark-colored  mass  floating  on  the  surface  of  the 
liquid.     Silicic  acid  sometimes  separates  as  a  white  gelatinous 
mass,  indicating  a  silicate.     [Compare  B.] 

In  any  case  filter,  wash  with  hot  water,  and  evaporate  the 
filtrate  to  dryness.  The  undissolved  residue,  if  any,  is  preserved 
for  D.  Dissolve  the  evaporated  filtrate  in  water  and  examine 
for  metals,  as  in  Part  III. 

D.  The  substance  is  insoluble  in  H20,  and  in  both  HC1  and  HN03. 
The  residue  from  C  is  next  treated  with  aqua  regia  under  a 


118  QUALITATIVE    ANALYSIS 

hood.  After  digesting  for  some  minutes,  the  mass  is  diluted 
with  an  equal  portion  of  water  and  filtered.  Any  residue  is 
preserved  for  E.  The  filtrate  is  evaporated  to  dry  ness,  dis- 
solved in  water,  and  the  metals  separated  as  before. 

E.  The  substance  is  insoluble  in  H20  and  in  all  acids.  The  residue 
from  D  may  consist  of  any  or  all  of  the  following  compounds, 
viz.  : 

Barium,  strontium,  and  calcium  sulfates, 

Lead  sulfate,  and  possibly  lead  chlorid, 

Silver  chlorid,  bromid,  iodid,  and  cyanid, 

Silicic  acid  and  many  silicates, 

Aluminum,  and  chromium  oxids, 

Calcium  fluorid, 

Sulfur,  and 

Carbon. 

In  addition  to  these,  a  few  rare  compounds  consisting  of  cer- 
tain aluminates,  phosphates,  arsenates,  and  oxids  are  insoluble 
in  acids,  and  so,  if  present,  would  be  found  here. 

Preliminary  tests  may  be  made  to  see  what  compounds  are 
present,  although  this  is  not  absolutely  essential.  It  is  better, 
however,  to  test  a  small  portion  before  each  step,  provided  the 
amount  of  material  is  not  too  small. 

(a)  Lead  and  silver  salts,  if  present,  must  first  be  removed. 
The  lead  salts  are  removed  by  heating  the  whole  residue  with 
a  concentrated  solution  of  ammonium  acetate,  which  dissolves 
them.     Filter,  wash,  and  test  one  portion  of  the  filtrate  with 
BaCl2  for  sulfuric  acid,  another  portion  with  H2S  for  lead,  and 
a  third  portion,  acidified  with  HNO3,  with  AgNO3  for  chlorin. 

(b)  If   a   residue    remains,    or   if   nothing    dissolves    in   the 
ammonium  acetate,  digest  the  residue  for  some  time  with  a 
solution  of  KCN,  at  a  gentle  heat,  repeating  the  operation,  if 
necessary,  until  all  the  silver  salts   are  dissolved.     (If  sulfur 
is  present,  this  operation  should  be  carried  on  in   the  cold.) 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS      119 

The  whole  is  now  filtered  and  the  precipitate  well  washed. 
If  the  filtrate  is  acidified  with  HNO3  (do  this  under  a  hood), 
the  silver  salt  is  precipitated  and  may  be  further  tested  for  silver 
by  filtering  and  fusing  a  portion  of  this  last  precipitate  on 
charcoal  with  Na2CO3,  which  will  give  a  bead  of  metallic  silver. 

(c)  The  residue  is  now  free  from  lead  and  silver  salts.  It  is 
next  heated  in  a  covered  porcelain  crucible  until  all  sulfur, 
if  present,  is  volatilized.  It  is  then  mixed  with  four  parts  of 
NaKCO3  and  one  part  of  KNO3  and  heated  in  a  platinum  crucible 
for  a  half  hour,  or  until  it  is  all  thoroughly  fused.  If  the  red- 
hot  crucible  is  placed  on  a  cold  metal  plate,  the  fused  mass  can 
generally  be  removed  in  a  cake  after  it  has  become  cold. 
Digest  the  fused  mass  in  hot  water  for  a  half  hour  and  then 
filter,  washing  the  residue  once  with  hot  water.  The  residue 
should  be  thoroughly  washed,  but  the  wash  water,  except  the 
first  washings,  may  be  thrown  away.  Preserve  the  residue 
for  examination  under  (d). 

The  filtrate  now  contains  the  acids  which  were  present  in 
the  insoluble  compounds,  together  with  those  bases  which  are 
soluble  in  the  alkaline  carbonates.  It  is  examined  as  follows. 

(tfj)  Take  a  small  portion  of  the  filtrate,  acidify  it  with  HC1, 
and  add  BaCl2.  A  white  precipitate  indicates  a  sulfate. 

(c2)  Take  another  portion,  acidify  it  with  H2SO4,  heat  it  to 
boiling,  and,  while  hot,  saturate  it  with  H2S.  A  yellow  pre- 
cipitate is  As2S3  and  indicates  an  arse-note. 

(CB)  Filter  off  the  yellow  arsenic  sulfid  from  (<?2),  add  to  the 
filtrate  enough  concentrated  HNO3  to  make  it  strongly  acid, 
and  then  some  ammonium  molybdate  solution.  A  yellow  pre- 
cipitate indicates  a  phosphate. 

(<?4)  To  another  portion  of  the  filtrate  add  HC1  to  decom- 
pose the  carbonates,  and  boil  until  all  CO2  has  been  driven  off. 
Make  the  solution  alkaline  with  NH4OH,  filter,  and  add  to  the 


120  QUALITATIVE  ANALYSIS 

filtrate  some  CaCl2.  If  a  precipitate  forms,  it  is  CaF2  and  indi- 
cates a  fluorid.  This  may  be  further  tested  as  in  Part  I, 
Hydrofluoric  Acid  5. 

(c5)  If  the  filtrate  from  (c)  is  yellow,  it  may  contain  chro- 
mates.  Acidify  a  portion  with  acetic  acid  and  add  lead  acetate. 
A  yellow  precipitate  indicates  a  chr ornate. 

(cQ)  The  remainder  of  the  filtrate  is  acidified  with  HC1  and 
evaporated  to  dryness.  It  is  best  to  repeat  this  operation, 
adding  more  HC1  to  insure  the  complete  decomposition  of  the 
silicic  acid.  The  dried  mass  is  then  treated  with  hot  water,  a 
few  drops  of  HC1  being  added  if  necessary. 

If  silicic  acid  was  present,  it  will  now  appear  as  an  insoluble 
precipitate  of  silicon  dioxid,  SiO2.  The  metals  which  were 
dissolved  by  the  carbonates  will  now  be  in  solution  as  chlorids. 
Filter  off  the  SiO2.  This  indicates  a  silicate,  and  may  be  tested 
as  in  Part  II,  V,  (h).  Test  the  filtrate  for  metals  as  in  Part  III. 

(d)  The  residue  from  (c)  is  now  washed  into  a  porcelain  dish, 
made  acid  with  HC1,  and  evaporated  to  dryness,  repeating  the 
operation,  as  directed  above,  to  insure  complete  decomposition 
of  any  silicic  acid  present.  Dissolve  the  dried  mass  in  water, 
add  a  few  drops  of  HC1,  filter  off  any  insoluble  silicon  dioxid, 
and  examine  the  filtrate  for  metals,  as  given  in  Part  III. 

F.  The  substance  is  a  silicate.  If  the  preliminary  examination 
(see  page  112)  shows  us  that  the  substance  is  a  silicate,  most  of 
the  preceding  steps  may  be  omitted,  or  at  least  modified. 

All  silicates  may  be  divided  into  two  general  classes,  viz.  : 

(a)  Silicates  which  are  completely  decomposed  by  the  mineral 
acids,  HC1,  H2S04,  or  HNO3. 

(b)  Silicates    that   are    not   decomposed,   or  which  are   only 
partially  decomposed,  by  acids. 

In  order  to  determine  to  which  of  these  two  classes  a  given 
silicate  belongs,  it  is  first  necessary  to  reduce  the  substance  to  a 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS       121 

very  fine  powder.  Digest  a  portion  of  this  powder  with  some 
HC1,  warming,  but  not  boiling,  the  liquid.  If  the  acid  decomposes 
the  silicate,  this  will  generally  be  indicated  by  a  change  in  the 
color  of  the  solution  and  the  presence  of  silicic  acid,  which 
appears  as  a  white,  flocculent,  gelatinous,  or  pulverulent  mass 
in  the  solution.  If  PIC1  fails  to  decompose  it,  test  another 
portion  with  H2SO4  (three  parts  acid  to  one  part  water).  The 
indications  are  the  same  as  with  HC1,  and  if  both  fail  to 
decompose  the  silicate,  it  belongs  to  the  second  class. 

(a)  Silicates  decomposed  by  acids.  The  finely  powdered  silicate 
is  mixed  with  a  little  water  to  a  thin  paste,  in  a  porcelain  dish, 
and  then  digested  with  HC1  at  a  temperature  near  the  boiling 
point,  until  it  is  completely  decomposed.  Evaporate  the  whole 
to  dry  ness,  with  occasional  stirring,  add  more  acid,  and  evaporate 
the  second  time.  Treat  the  dried  mass  with  hot  water,  adding 
a  few  drops  of  HC1,  and  filter  off  the  insoluble  silica,  SiO2. 

The  metals  will  now  be  in  solution  as  chlorids  and  may  be 
separated  as  in  Part  III. 

If  lead  or  silver  is  found  in  the  silicate,  it  is  better  to  use 
HNO3,  the  operation  being  otherwise  conducted  as  above. 
Sulfuric  acid  may  be  used  if  an  aluminum  silicate  is  to  be 
decomposed. 

Silicates  not  infrequently  contain  small  quantities  of  titanium 
in  the  form  of  TiO2,  and  replacing  some  of  the  SiO2.  If  the 
solution,  after  filtering  off  the  silica,  is  diluted  and  then  boiled 
for  a  long  time,  the  TiO2  is  precipitated.  (Titanium  is  a  rare 
element,  which  need  not  be  looked  for  in  most  qualitative 
analyses.) 

(6)  Silicates  not  decomposed  by  acids.  Take  about  a  gram  of  the 
finely  powdered  silicate  and  mix  it  with  about  four  grams  of 
NaKCO3.  (A  mixture  of  equal  parts  of  Na2CO3  and  K2C3O 
will  answer.)  Place  the  mixture  in  a  platinum  crucible  and 
heat  for  half  an  hour,  or  until  it  is  all  in  a  state  of  complete 


122  QUALITATIVE  ANALYSIS 

fusion.  If  the  crucible  is  allowed  to  cool  on  a  metal  plate,  the 
fused  mass  can  usually  be  removed  in  a  lump. 

Place  this  fused  mass,  or  the  crucible  and  its  contents,  if 
the  contents  cannot  be  easily  removed,  in  a  porcelain  dish, 
cover  with  warm  water,  and  add  HC1  sufficient  to  decompose 
the  carbonates.  The  whole  is  now  evaporated  to  dryness  and 
treated  as  in  the  case  of  silicates  decomposed  by  acids. 

Of  course  after  the  addition  of  both  sodium  and  potassium 
carbonates,  this  portion  cannot  be  examined  for  alkali  metals. 
If  these  are  present,  they  may  be  detected  as  follows. 

Take  another  portion  of  the  finely  powdered  silicate,  mix 
with  it  one  part  of  ammonium  chlorid  and  eight  parts  of  pre- 
cipitated calcium  carbonate.  Heat  this  mixture  to  moderate 
redness  for  half  an  hour  in  a  covered  platinum  crucible.  Place 
the  crucible  and  its  contents  in  a  porcelain  dish  with  water, 
boil  until  the  whole  mass  is  completely  disintegrated,  and 
filter.  Add  to  the  filtrate  NH4OH  until  it  is  alkaline,  then 
ammonium  carbonate  to  slight  excess.  See  that  the  mixture 
now  smells  strongly  of  ammonia,  allow  it  to  settle,  and  filter. 

The  last  filtrate  will  now  contain  the  alkali  metals  and  of 
course  ammonium  salts.  Evaporate  this  to  dryness  and  ignite 
gently  until  all  the  ammonium  compounds  have  been  decom- 
posed. Dissolve  what  remains  in  water,  filter  if  necessary,  and 
examine  the  solution  as  in  Part  III,  Group  6. 

G.  Cyanids  are  present.  The  cyanids,  and  more  especially  the 
compound  cyanids,  if  present,  are  not  completely  dissolved  or 
decomposed  when  treated  with  HC1,  as  in  division  B.  They 
will  generally  be  detected  by  the  methods  already  given,  but 
a  much  more  satisfactory  method  of  treatment  when  they  are 
present  is  the  following. 

A  small  portion  of  the  residue  from  A  is  tested  with  HC1, 
which  will  usually  decompose  some  of  the  cyanids  so  as  to  liber- 
ate HCN.  If  cyanids  are  found,  take  the  whole  residue  from 


SYSTEMATIC  EXAMINATION  OF  COMPLEX  SOLIDS      123 

A  and  boil  it  with  a  little  concentrated  NaOH  solution  for  a' 
short  time,  add  some  concentrated  solution  of  Na2CO3  and  boil 
again  for  a  few  minutes.  Dilute  with  hot  water,  filter,  and 
wash  the  residue,  if  any,  with  hot  water.  The  residue  will  now 
be  free  from  all  cyanids,  except  silver  cyanid  (which  will  be 
found  in  E),  and  is  treated  as  given  under  B. 

(a)  A  small  portion  of  the  strongly  alkaline  filtrate  is  treated 
with  H2S.     (It  is  best  to  use  a  solution  of  H2S ;  or,  if  the  gas 
itself  is  used,  not  to  completely  saturate  the  solution,  as  some 
of  the  metals,  such  as  aluminum,  etc.,  may  be  precipitated,  since 
NaOH  is  present.)     If  no  precipitate  forms,  pass  on  to  (b).     If 
a  precipitate  forms,  take  about  one  half  of  the  alkaline  filtrate 
and  add  Na2S   [(NH4)2S  may  be  used],  avoiding  any  unneces- 
sary excess.     Warm  the  whole  without  boiling,  filter,  and  wash 
with  hot  water,  saving  the  filtrate  for  (b). 

The  precipitate,  which  will  contain  the  metals  belonging  to 
Groups  1,  2  (Subdivision  A),  3,  and  4,  is  treated  with  dilute 
HNO3.  Filter  off  any  insoluble  HgS  or  PbSO4  which  may  be 
present  and  examine  for  metals,  as  given  in  Part  III. 

(b)  The  filtrate  from  (a)  or  a  portion  of  the  alkaline  filtrate, 
if  no  precipitate  was  formed  by  the  Na2S,  is  now  made  acid  with 
dilute  H2SO4  and  saturated  with  H2S.     If  no  precipitate  forms, 
pass  on  to  (c).     If  there  is  a  precipitate,  it  belongs  to  Group  2, 
Subdivision  B.     Examine  as  given  on  page  84. 

(c)  The  filtrate  from  (b)  still  contains  those  metals  which  form 
the  double  cyanogen  compounds  (ferro-  and  ferri-cyanids,  cobalti- 
cyanids,  etc.)  and  aluminum.     A  portion  of  the  solution  may  be 
tested  for  the  acid  radicals,  as  given  in  Part  III. 

Evaporate  another  portion  of  this  filtrate  nearly  to  dryness 
after  acidifying  strongly  with  H2SO4,  and  then  heat  (under  a 
hood)  until  most  of  the  excess  of  H2SO4  has  been  driven  off. 
Dissolve  the  residue  in  water  and  examine  the  solution  for 
metals  of  the  third  and  fourth  groups,  as  given  in  Part  III. 


124  QUALITATIVE   ANALYSIS 

Since  the  ferricyanids  are  reduced  to  ferrocyanids  by  H2S, 
and  since  H2SO4  is  used  in  the  tests  for  the  other  acids  and 
metals,  we  must  test  for  sulfuric  and  ferricyanic  acids  in  the 
remaining  portion  of  the  original  alkaline  solution.  This  may 
be  done  in  the  usual  way  after  acidifying  with  HC1  or  HNO3. 
[See  the  corresponding  acids  in  Part  I.] 


APPENDIX  A 


Names,  Symbols,  and  Atomic  Weights  of  the  Elements 

The  non-metallic  elements  are  printed  in  small  capitals. 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 
O  =  1G 

NAME 

SYMBOL 

ATOMIC 
WEIGHT 
O  =  1G 

Aluminum  

Al. 

27.1 

Neodymium  

Nd. 

143.6 

Antimony 

Sb. 

120.2 

NEON 

Ne. 

20. 

ARGON 

A. 

39.9 

Nickel 

Ni. 

58.7 

ARSENIC 

As. 

75. 

NITROGEN 

N. 

14. 

Barium 

Ba. 

137.4 

Osmium 

Os. 

191. 

Bismuth  
BORON  
BROMIN  
Cadmium  

Bi. 
B. 

Br. 

Cd. 

208.5 
11. 
80. 
112.4 

OXYGEN  
Palladium  
PHOSPHORUS  
Platinum  

0. 

Pd. 
p 

Pt. 

16. 
106.5 
31. 
194.8 

Caesium 

Cs. 

133. 

Potassium 

K. 

39  1 

Calcium 

Ca. 

40.1 

Praseodymium 

Pr. 

140  5 

CARBON 

C. 

12. 

Radium 

Rd. 

225. 

Cerium 

Ce. 

140. 

Rhodium 

Rh. 

103. 

CHLORIN 

01. 

35.4 

Rubidium 

Rb. 

85.4 

Chromium 

Cr. 

52.1 

Ruthenium 

Ru. 

101.7 

Cobalt  
Columbium  
Copper  

Co. 
Cb. 
Cu. 

59. 
94. 
63.6 

Samarium  
Scandium  
SELENIUM  

Sm. 

Sc. 
Se. 

150. 
44.1 
79.2 

Erbium 

Er. 

166. 

SILICON 

Si. 

28  4 

FLUORIN 

F. 

19. 

Silver 

As. 

107  9 

Gadolinium 

Gd. 

156. 

Sodium 

Na. 

23 

Gallium.  
Germanium  
Glucinum  
Gold  
HELIUM  

Ga. 
Ge. 
Gl. 
Au. 
He. 

70. 
72.5 
9.1 
197.2 
4. 

Strontium  
SULFUR  
Tantalum  
TELLURIUM  
Terbium          

Sr. 
S. 
Ta. 
Te. 
Tb. 

87.6 
32. 
183. 
127.6 
160. 

HYDROGEN  
Indium 

H. 
In. 

1.0075 
114. 

Thallium  
Thorium 

Tl. 
Th. 

204.1 
232  5 

IODIN 

I. 

126.8 

Thulium 

Tu. 

171. 

Indium 

Ir. 

193. 

Tin 

Sn. 

119 

Iron  
KRYPTON  
Lanthanum 

Fe. 
Kr. 
La. 

55.9 
81.8 
138.9 

Titanium  
Tungsten  
Uranium 

Ti. 
W. 

u 

48.1 
184. 
238  5 

Lead 

Pb. 

206.9 

Vanadium 

v 

51  2 

Lithium  
Magnesium  
Manganese  
Mercury  
Molybdenum  

Li. 
Mg. 
Mn. 
Hg. 
Mo. 

7. 
24.3 
55. 
200. 
96. 

XENON  
Ytterbium  
Yttrium  
Zinc  
Zirconium  

Xe. 

Yb. 
Yt. 
Zn. 
Zr. 

128. 
173. 
89. 
65.4 
90.4 

125 


APPENDIX  B 
Names  and  Formulas  of  Reagents  and  Solutions  used  in  this  Work 

In  the  formulas  given  below,  the  water  of  crystallization  has  been  omitted. 

Acetic  Acid H(C2H302)  or  HAc 

Acid  Sodium  Phosphate Na2HP04 

Acid  Sodium  Tartrate         NaH(C4H406) 

Alcohol C2H5OH 

Aluminum  Sulfate A12(S04)3 

Ammonium  Acetate NH4(C2H302) 

Ammonium  Carbonate (NH4)2C03 

Ammonium  Chlorid NH4C1 

Ammonium  Hydroxid NH4OH 

Ammonium  Molybdate (NH4)2Mo04 

Ammonium  Oxalate (NH4)2C204  or  (NH4)2  Ox 

Ammonium  Sodium  Hydrogen  Phosphate     .     .     .       NH4NaHP04 

Ammonium  Sulfid (NH4)2S 

Ammonium  Sulfid  (yellow) (NH4)2SX 

Ammonium  Tartrate (NH4)2(C4H40C) 

Antimony  Chlorid SbCl3 

AquaRegia [3  HC1  +  HN03] 

Arsenious  Oxid As203  or  As4O6 

Barium  Carbonate BaC03 

Barium  Chlorid BaCl2 

Barium  Hydroxid Ba(OH)2 

Baryta  Water [Ba(OH)2  +  (H2O)X] 

Bismuth  Nitrate Bi(N03)3 

Borax Na2B407 

Bromin  Water       .     .     '. [Br  +  (H,0)  J 

Cadmium  Nitrate Cd(N03)2 

Calcium  Chlorid CaCl2 

Calcium  Hydroxid Ca(OH)2 

Carbon  Disulfid .     .    CS2 

126 


APPENDIX  127 

Chlorin  Water [Cl  +  (H20)x] 

Chromium  Sulfate Cr2(S04)3 

Cobalt  Nitrate Co(N03)2 

Copper  Sulfate CuS04 

Ether (C2H5)20 

Ferric  Chlorid        FeCl3 

Ferrous  Chlorid FeCl2 

Ferrous  Sulfate FeS04 

Hydrochloric  Acid HC1 

Hydrofluosilicic  Acid H2SiF6 

Hydrogen  Sulfid H2S 

Lead  Acetate Pb(C2H302)2  or  PbAc, 

Lead  Nitrate Pb(N03)2 

Lime  Water [Ca(OH)2  +  (H20)J 

Magnesium  Sulfate MgS04 

Manganese  Dioxid Mn02 

Manganese  Sulfate MnS04 

Mercuric  Chlorid HgCl2 

Mercuric  Nitrate Hg(N03)2 

Mercurous  Nitrate HgN03 

Microcosmic  Salt NH4NaHP04 

Nessler's  Eeagent [HgI2(KI)2  +  (KOH)X] 

Nickel  Nitrate Ni(N03)2 

Nitric  Acid       HN03 

Oxalic  Acid H2C204  or  H2Ox 

Platinum  Chlorid PtCl4 

Potassium  Arsenate K3As04 

Potassium  Arsenite K3AsG3 

Potassium  Bichromate K2Cr207 

Potassium  Bromid KBr 

Potassium  Chlorate KC103 

Potassium  Chromate K2Cr04 

Potassium  Cyanid KCN 

Potassium  Ferricyanid K3Fe(CN)G 

Potassium  Ferrocyanid K4Fe(CN)6 

Potassium  Fluorid KF 

Potassium  lodid  KI 


128  QUALITATIVE   ANALYSIS 

Potassium  Nitrate  (Saltpeter) KN03 

Potassium  Nitrite KN02 

Potassium  Pyroantimonate  (Acid) K2H2Sb207 

Potassium  Sulfocyanate KSCN 

Potassium  Thiocyanate        KSCN 

Rochelle  Salt NaK(C4H4Oc) 

Saltpeter KN03 

Silver  Nitrate        AgN03 

Sodium  Acetate Na(C2H302)  or  NaAc 

Sodium  Arsenate NasAs04 

Sodium  Carbonate Na2C03 

Sodium  Hydroxid NaOH 

Sodium  Hypochlorite NaCIO 

Sodium  Metasilicate Na2Si03 

Sodium  Phosphate  (Acid)        Na2HP04 

Sodium  Pyroborate  (Borax) , Na2B407 

Sodium  Silicate Na4Si03 

Sodium  Sulfate Na2S04 

Sodium  Sulfite Na2S03 

Sodium  Tartrate  (Acid) NaH(C4H406)  or  NaH  Tr" 

Sodium  Thiosulfate Na2S203 

Stannic  Chloric! SnCl4 

Stannous  Chlorid        SnCl2 

Strontium  Chlorid SrCl2 

Sulfuric  Acid H2S04 

Sulfurous  Acid H2SO3 

Tartaric  Acid H2(C4H406)  or  H2Tr 

Zinc  Sulfate  ZnSO4 


APPENDIX   C 

Preparation  of  Reagents  and  Solutions  for  Analysis 

The  strength  of  the  solution,  both  of  the  reagent  and  the  sub- 
stance for  analysis,  is  quite  an  important  factor  in  the  analysis. 
If  the  solutions  are  too  strong,  the  bulk  of  the  precipitate  is  often 
so  great  as  to  make  the  mass  almost  solid.  This  makes  the  pre- 
cipitate difficult  to  handle  and  adds  unnecessarily  to  the  expenses 
of  the  laboratory.  On  the  other  hand,  if  the  solutions  are  too 
dilute,  some  ingredient  may  fail  of  precipitation  and  so  be  missed 
in  the  analysis.  Furthermore,  solutions  made  up  in  a  haphazard 
way  will  cause  much  annoyance  in  precipitation,  especially  in  the 
more  advanced  portions  of  the  work;  it  will  be  found  to  be  much 
more  convenient  if  the  solutions  are  chemically  equivalent,  or  of 
such  relative  strength  that  chemically  equivalent  amounts  may  be 
used.  All  solutions  and  dilutions  are  made  with  water,  unless 
otherwise  stated. 

The  normal  solution  of  volumetric  analysis  is  the  best  standard  of 
strength.  A  normal  solution  is  a  solution  so  prepared  that  one 
liter  shall  contain  the  hydrogen  equivalent  of  the  active  reagent, 
weighed  in  grams.  In  the  case  of  substances  in  which  the  active 
reagent  is  univalent,  the  hydrogen  equivalent  is  the  same  as  the 
molecular  weight.  When  the  active  reagent  is  bivalent  or  trivalent, 
the  hydrogen  equivalent  is  one  half  or  one  third  of  the  molecular 
weight  respectively.  Thus  the  hydrogen  equivalent  of  NaOH  is  the 
molecular  weight  of  NaOH,  or  40,  and  one  liter  of  water,  contain- 
ing 40  grains  of  NaOH,  is  therefore  a  normal  solution.  Again,  the 
hydrogen  equivalent  of  H2S04  is  one  half  of  its  molecular  weight, 
since  it  contains  two  replaceable  hydrogen  atoms.  The  molecular 
weight  of  H2S04  is  98,  and  one  half  of  this  number,  or  49  grams, 
of  H2S04  in  a  liter  makes  a  normal  solution.  In  a  normal  solution 
of  a  salt  the  water  of  crystallization  must  also  be  taken  into  account. 
Thus  the  compound  called  microcosmic  salt  is  NH4NaHP04  +  4  H20, 

129 


130  QUALITATIVE   ANALYSIS 

and  its  molecular  weight  is  209,  which  is  three  times  its  hydrogen 
equivalent,  phosphoric  acid  being  tribasic,  so  that  69.6  grams  of 
this  salt  in  a  liter  makes  a  normal  solution.  In  preparing  these 
solutions  it  is  not  necessary  to  more  than  closely  approximate 
these  weights. 

Normal  solutions  are  marked  N,  half-normal  solutions  |,  and 
double-normal  solutions  2  N.  For  most  operations  one  half  normal, 
or  even  one  fourth  normal,  solutions  will  be  of  sufficient  strength. 
Some  reagents  may  be  used  more  dilute  than  this.  Thus  AgN03 
may  be  used  one  fifth  or  one  tenth  normal  for  most  reactions. 

The  concentrated  acids  have  approximately  the  following  strength, 
viz.:  H2SO4,  specific  gravity  1.84,  about  36  N ;  HN03,  specific 
gravity  1.42,  about  16  N;  HC1,  specific  gravity  1.20,  about  12  N ; 
acetic  acid,  30  per  cent,  about  5  N.  Concentrated  ammonia,  specific 
gravity  0.90,  has  a  strength  of  about  20  N. 

These  acids,  and  ammonia,  are  often  used  for  other  purposes  than 
precipitation.  It  will  be  better,  therefore,  to  make  these  somewhat 
stronger  than  the  other  reagents.  About  four  times  normal  (4  N) 
will  be  found  a  convenient  strength.  H2S04  diluted  1 :  8  is  about 
4  N ;  HN03  diluted  1 : 3  is  about  4  N ;  HC1  diluted  1 : 2  is  about 
4  N ;  ammonia  diluted  1  : 4  is  about  4  N.  It  will  also  be  found 
convenient  to  have  solutions  of  other  reagents  about  4  N  in 
strength. 

A  few  reagents  are  so  difficultly  soluble  in  water  that  only  very 
dilute  solutions  can  be  obtained.  A  saturated  solution  of  bromin  in 
water  is  about  ^ ;  chlorin  water  is  about  N ;  barium  hydroxid 
(baryta  water)  about  §;  calcium  hydroxid  (lime-water)  about  ^; 
calcium  sulfate  about  ^. 

The  strength  of  other  solutions  may  be  varied  according  to  the 
purpose  for  which  they  are  employed. 


INDEX 


PAGE 

Acetates,  test  for 55 

Acetic  acid,  reactions  for  .  45,  65,  67 

separation  of 110 

Acid  radicals,  classification  of  .  103 

examination  for       .     .    75,  101 

solution  for 102 

reactions  for 30 

Acids,  Group  1 104 

Group  2 107 

Group  3 110 

Alkali  metals,  the 27 

Alloys,  examination  of  ...  113 
Aluminum,  reactions  for  .  .  .  13 

separation  of  ....  86 

Ammonia,  test  for 51 

Ammonium,  reactions  for  .  29,  53 

separation  of 98 

Antimony,  reactions  for  ...  10 

separation  of 84 

Antimony  sulfid,  test  for  .  .  .  53 
Arsenic,  reactions  for  .  .  10,  53,  68 

separation  of 84 

Arsenic  acid,  reactions  for  .  .  42 

separation  of 105 

Arsenic  sulfid,  test  for  ....  54 
Arsenious  acid,  reactions  for  .  41 

separation  of 105 

Atomic  weights,  table  of  .  .  .  125 


Barium,  reactions  for   . 

separation  of 
Bismuth,  reactions  for  . 

separation  of  .     . 
Blowpipe 


24 
96 

7 

82 
49 


PAGE 

Blowpipe  analysis 48 

Blowpipe  flame 49 

Borax  bead,  test  with  ....  60 

Boric  acid,  reactions  for    .     .     43,  67 

separation  of       ....  93 

Bromin,  test  for 52 

Bunsen  lamp 49 

Cadmium,  reactions  for     ...  9 

separation  of       ....  82 

Calcium,  reactions  for  ....  26 

separation  of 96 

Carbon  dioxid,  test  for     .     44,  51,  63 
Carbon  monoxid,  test  for  .     .     51,  64 

Carbonic  acid,  reactions  for  .     .  43 

separation  of 105 

Charcoal,  examination  on      .     .  55 

with  sodium  carbonate     .  58 

Chloric  acid,  reactions  for      .     .  35 

separation  of 110 

Chlorin,  test  for 52 

Chlorin  peroxid,  test  for    ...  65 

Chromic  acid,  reactions  for    .     .  38 

separation  of       ....  105 

Chromium,  reactions  for  ...  14 

separation  of       ....  86 

Citric  acid,  separation  of  .     .     .  106 

Closed  tube,  examination  in  .  50 

Cobalt,  reactions  for     ....  19 

separation  of      ....  93 

Coloration  of  borax  bead  ...  60 

Coloration  of  flame 59 

Complex   solids,   examination 

of  .  Ill 


131 


132 


INDEX 


Copper,  reactions  for    .     . 
separation  of       .     . 
Cyanids,  special  analysis  of 
Cyanogen  gas,  test  for 

Elements,  table  of    ... 


Formates,  test  for 


PAGE 

8 

82 

122 

52 

125 

55 


Lead,  reactions  for 
separation  of 


PAGE 
1 

78,  82 


Heat  alone,  effects  of    ....       50 
Hydriodic  acid,  reactions  for  32,  64,  66 
separation  of       ....     109 
Hydrobromic     acid,    reactions 

for 31,  64,  66 

separation  of       ....     109 
Hydrochloric    acid,    reactions 

for 31,  64,  66 

separation  of       ....     109 
Hydrocyanic  acid,  reactions  for      33 
separation  of  ...       108,  122 
Hydroferricyanic  acid,  reactions 

for 34 

separation  of       .     .       108,  122 
Hydroferrocyanic  acid,  reactions 

for 34 

separation  of       .     .       108,  122 

Hydrofluoric  acid,  reactions  for  32,  64 

separation  of       ....       93 

Hydrogen,  test  for 65 

Hydrogen  sulfid,  reactions  for 

36,  52,  64 

separation  of       ....     107 
Hydrosulfuric  acid.   (See  Hydro- 
gen sulfid.) 
Hypochlorous  acid,  reactions  for      35 

lodin,  test  for 52,  53 

Iron  (ferric),  reactions  for     .     .       16 

separation  of       ....       86 

Iron  (ferrous),  reactions  for  .     .       15 


Magnesium,  reactions  for  ...  23 

separation  of       ....  98 

Manganese,  reactions  for  .     .     .  20 

separation  of       ....  93 

Mercuric  compounds,  test  for     53,  54 

Mercurous  chlorid,  test  for    .     .  53 

Mercury  (mercuric) ,  reactions  for  6 

separation  of       ....  82 

Mercury  (mercurous),  reactions 

for 4 

separation  of       ....  78 

Metals  and  alloys,  examination  of  113 
Metals,  separation  of.  (See  Mixed 
compounds.) 

Mixed  compounds 76 

Group  1 78 

Group  2 80 

Group  2,  subdivision  A    .  82 

Group  2,  subdivision  B     .  84 

Group  3 86 

Group  4     ......  93 

Group  5 96 

Group  6 98 

Nickel,  reactions  for  ....  17 

separation  of  ....  93 

Nitric  acid,  reactions  for  .  .  40,  65 

separation  of  ....  110 

Nitrogen  oxids,  tests  for    .     .      52,  65 

Nitrous  acid,  reactions  for  .  .  39 

separation  of  ....  110 

Organic  gases,  test  for  .     ...  52 

Oxalic  acid,  reactions  for  .     .     .  46 

separation  of       .     .    ".     91,  92 

Oxidizing  flame 49 

Oxygen,  test  for 51,  63 


INDEX 


133 


PAGE 

Peroxids,  test  for 68 

Phosphates,  etc.,  absent    ...  88 

present 90 

Phosphates,  test  for      ....  G8 
Phosphoric  acid,  reactions  for     41,  68 

separation  of       ...     00,  92 
Platinum  foil,  substances  fused 

on 62 

Potassium,  reactions  for    ...  27 

separation  of       ....  98 

Reagents,  table  of,  with  formulas  126 

Reducing  flame 49 

Silicates,  analysis  of  ....  120 

decomposed  by  acids  .  .  121 

not  decomposed  by  acids  .  121 

Silicic  acid,  reactions  for  ...  44 

separation  of  .  .  .  91,  120 

Silver,  reactions  for  ....  3 

separation  of  ....  78 

Simple  compounds 69 

Group  1 69 

Group  2  ......  70 

Group  3 72 

Group  4 73 

Group  5 73 

Group  6 74 

Sodium,  reactions  for  ....  28 

separation  of  ....  98 


PAGE 

,  Solutions,  directions  for  making  129 
Special  tests  in  blowpipe  analysis  66 
Strontium,  reactions  for  ...  25 

separation  of 96 

Sulfates,  test  for 67 

Sulfocyanic  acid,  reactions  for  34,  67 

separation  of  ....  107 

Sulfur,  test  for 54 

Sulfur  compounds,  test  for  .  59,  66 
Sulfur  dioxid,  test  for  .  .  52,  65,  67 
Sulfuric  acid,  reactions  for  .  38,  67 

separation  of       ....     104 

substances  heated  in  .  .  62 
Sulfurous  acid,  reactions  for  .  37,  67 

separation  of       ....     104 

Tartaric  acid,  action  as  reagent   13,  16 

reactions  for 46 

separation  of       ....     106 

Tartrates,  test  for 55 

Thiocyanic  acid.  (See  Sulfocyanic 

acid. ) 

Thiosulfuric  acid,  reactions  for  36,  67 

Tin  (stannic),  reactions  for    .     .       12 

separation  of       ....       84 

Tin  (stannous),  reactions  for      .       11 

separation  of       ....       84 


Zinc,  reactions  for 
separation  of 


22 
93 


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IN  ONE  VOLUME.  I2mo.  Cloth.  564  pages.  Illustrated,  partly  in  color. 
List  price,  $2.00;  mailing  price,  $2.15. 

IN  TWO  VOLUMES.  Volume  I.  Elements  of  Physiology.  lamo.  Cloth. 
321  pages.  Illustrated.  List  price,  ;  mailing  price,  .  Volume  II. 
Hygiene  and  Sanitation.  I2mo.  Cloth.  271  pages.  Illustrated.  List 
price,  ;  mailing  price, 


THE  recent  advances  in  our  knowledge  of  personal 
and  public  hygiene  and  of  sanitation  have  made 
indispensable  for  the  use  of  progressive  educators  a 
practical  and  authoritative  text-book  of  what  may  be  called 
"  the  new  physiology  and  hygiene."  The  present  work  has 
been  prepared  in  recognition  of  this  need.  Professor  Hough 
has  been  engaged  in  the  study  and  teaching  of  physiology 
and  personal  hygiene  for  over  ten  years.  Professor  Sedg- 
wick  has  had  a  large  share  in  the  advancement  of  the  study 
of  public  health  and  sanitation  in  America.  They  have 
made  the  keynote  of  this  work  the  right  conduct  of  the 
physical  life,  and  to  this  end  everything  else  is  subordi- 
nated. Anatomy  and  histology  are  only  briefly  outlined, 
while  special  chapters  are  devoted  to  practical  matters  of 
hygiene  and  sanitation. 

The  authors  believe  that  the  matter  and  method  of  this 
new  text-book  will  go  far  toward  rescuing  physiology  and 
hygiene  from  the  neglected  and  too  often  despised  place 
which  they  now  occupy  in  schools,  and  promoting  them  to 
one  of  practical  usefulness  and  corresponding  esteem. 

The  book  is  intended  for  high  and  normal  schools,  and 
for  short  courses  in  colleges. 


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A  TEXT-BOOK  IN  GENERAL 
ZOOLOGY 

By  HENRY  R.  LINVILLE,  Head  of  the  Department  of  Biology,  De  Witt  Clinton 
High  School,  New  York  City,  and  HENRY  A.  KELLY,  Director  of  the  Depart- 
ment of  Biology  and  Nature  Study,  Ethical  Culture  School,  New  York  City 


462  pages.      Illustrated.      List  price,  $1.50  ;  mailing  price,  $1.70 


A  exposition  of  the  science  of  zoology,  presented  without  the  inter- 
polation of  a  laboratory  guide. 
Four  years  spent  in  careful  examination  of  the  original  sources 
have  resulted  in  a  book  filled  with  valuable  material.     The  authors, 
through   their  extended  service   as  teachers  of  biology  in  secondary 
schools,  are  well  equipped  for  the  task  of  writing  a  text-book  designed, 
as  this  one  is,  chiefly  for  high-school  use,  although  intended  also  to  be 
available  for  elementary  college  classes. 

The  structure,  the  physiology,  and  the  natural  history  of  selected 
types  of  animals  are  described  with  accuracy  and  in  language  easily 
understood  by  young  students. 

The  treatment  of  the  subject  is  broad,  and  the  inductive  method  is 
employed  so  far  as  each  class  and  phylum  of  invertebrate  animals  is 
concerned.  The  definition  of  a  group  is  not  given  until  the  student's 
conception  of  the  group  characters  has  grown  to  the  point  where  the 
definition  forms  the  fitting  end  to  the  logical  process  involved  in  the 
exposition. 

The  Insecta  are  discussed  in  the  first  chapters,  and  after  the  remainder 
of  the  Arthropoda  are  described,  the  other  invertebrate  phyla  follow  in 
a  descending  series,  ending  with  the  Protozoa.  Then,  beginning  with 
the  fishes,  the  order  ascends  to  the  mammals  and  closes  with  man. 

A  large  portion  of  the  book  is  devoted  to  the  insects  and  vertebrates, 
because  young  students  are  more  familiar  with  these  groups  and  so  take 
greater  interest  in  them.  The  less  known,  however,  are  treated  with  care, 
the  different  features  of  morphology,  physiology,  and  relation  to  envi- 
ronment being  maintained  in  good  measure  throughout. 

The  book  includes  two  hundred  and  thirty-three  illustrations. 


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A   FIRST   COURSE   IN    PHYSICS 

By  ROBERT  A.  MILLIKAN,  Assistant  Professor  of  Physics  in 

the    University   of  Chicago,    and    HENRY    G.    GALE, 

Instructor  in  Physics  in  the  University  of  Chicago 

iimo.     Cloth.     488  pages.     Illustrated.     List  price,  $1.25  ;  mailing  price,  $1.40 


A  LIST  OF  LABORATORY  EXPERIMENTS  IN  PHYSICS 

FOR  SECOND  ART  SCHOOLS 

By  ROBERT  A.   MILLIKAN  and  HENRY  G.   GALE 
List  price,  40  cents ;  mailing  price,  45  cents 


THIS  one-year  course  in  physics  has  grown  out  of  the 
experience  of  the  authors  in  developing  the  work  in 
physics  at  the  School  of  Education  of  the  University  of 
Chicago,  and  in  dealing  with  the  physics  instruction  in  affiliated 
high  schools  and  academies. 

The  book  is  a  simple,  objective  presentation  of  the  subject  as 
opposed  to  a  formal  and  mathematical  one.  It  is  intended  for 
the  third-year'  high-school  pupils  and  is  therefore  adapted  in  style 
and  method  of  treatment  to  the  needs  of  students  between  the 
ages  of  fifteen  and  eighteen.  It  especially  emphasizes  the  his- 
torical and  practical  aspects  of  the  subject  and  connects  the  study 
very  intimately  with  facts  of  daily  observation  and  experience. 

The  authors  have  made  a  careful  distinction  between  the  class 
of  experiments  which  are  essentially  laboratory  problems  and 
those  which  belong  more  properly  to  the  class  room  and  the 
lecture  table.  The  former  are  grouped  into  a  Laboratory  Manual 
which  is  designed  for  use  in  connection  with  the  text.  The  two 
books  are  not,  however,  organically  connected,  each  being  com- 
plete in  itself. 

All  the  experiments  included  in  the  work  have  been  carefully 
chosen  with  reference  to  their  usefulness  as  effective  class-room 
demonstrations. 


GINN  &   COMPANY  PUBLISHERS 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
BERKELEY 

Return  to  desk  from  which  borrowed. 
This  book  is  DUE  on  the  last  date  stamped  below. 


MAY    5  1948 


(9Mfl 


JUN  3 
JUN   3   1948 


t*& 


SCAugSlLf 


18May'62Xi, 


LD  21-100m-9,'47(A5702sl6)47 


MAY  2  9 1956  LO 

*       3Mr'59Wj 
REC'D  LD 

MAR  3     1959 


REC'D  LD 

JUL  1 5  1962 


APR  2  *  1963 


IMS  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


D 


D 


